Chemistry Form Four Best Notes Free

18.0.0 ACIDS, BASES AND SALTS

 

(25 LESSONS)

 

 

 

A.ACIDS AND BASES

At a school laboratory:

(i)An acid may be defined as a substance that turn litmus red.

 

(ii)A base may be defined as a substance that turn litmus blue.

 

Litmus is a lichen found mainly in West Africa. It changes its colour depending on whether the solution it is in, is basic/alkaline or acidic.It is thus able to identify/show whether

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  1. An acid is a substance that dissolves in water to form H+/H3O+ as the only positive ion/cation. This is called the Arrhenius definition of an acid. From this definition, an acid dissociate/ionize in water releasing H+ thus:

 

 

HCl(aq) -> H+ (aq)
HNO3(aq) -> H+ (aq)
CH3COOH(aq) -> H+ (aq)
H2SO4(aq) -> 2H+ (aq)
H2CO3(aq) -> 2H+ (aq)
H3PO4(aq) -> 3H+ (aq)

 

 

  • Cl (aq)
  • NO3 (aq)
  • CH3COO(aq)

 

  • SO42-(aq)

 

  • CO32-(aq)

 

  • PO43-(aq)

 

 

 

2.A base is a substance which dissolves in water to form OH as the only negatively charged ion/anion.

This is called Arrhenius definition of a base.

From this definition, a base dissociate/ionize in water releasing OH thus:

 

KOH(aq) -> K+(aq) + OH(aq)
NaOH(aq) -> Na+(aq) + OH(aq)
NH4OH(aq) -> NH4 +(aq) + OH(aq)
Ca(OH)2(aq) -> Ca2+(aq) + 2OH(aq)
Mg(OH)2(aq) -> Mg2+(aq) + 2OH(aq)

 

  1. An acid is a proton donor. A base is a proton acceptor.

This is called Bronsted-Lowry definition of acids and bases.

 

From this definition, an acid donates H+ .

H+ has no electrons and neutrons .It contains only a proton.

Examples

  1. From the equation:

 

HCl(aq)           +              H2O(l)            ===     H3O+(aq)       +            Cl (aq)

 

(a)(i)For the forward reaction from left to right, H2O gains a proton to form H3O+ and thus H2O is a proton acceptor .It is a Bronsted-Lowry base

 

  • For the backward reaction from right to left, H3O+ donates a proton to form H2O and thus H3O+ is an opposite proton donor. It is a Bronsted-

 

Lowry conjugate acid

 

(b)(i)For the forward reaction from left to right, HCl donates a proton to form Cl and thus HCl is a proton donor .

It is a Bronsted-Lowry acid

 

  • For the backward reaction from right to left, Cl gains a proton to form HCl and thus Cl is an opposite proton acceptor.

 

It is a Bronsted-Lowry conjugate base.

 

Every base /acid from Bronsted-Lowry definition thus must have a conjugate product/reactant.

 

II. From the equation:
HCl(aq) +(aq)    + Cl (aq)
+ NH3(aq) =  NH4
==
(a)(i)For the forward reaction from left to right, NH3 gains a proton to form
NH4 + and thus NH3 is a proton acceptor .

It is a Bronsted-Lowry base

 

  • For the backward reaction from right to left, NH4+ donates a proton to form NH3 and thus NH4+ is an opposite proton donor.

 

It is a Bronsted-Lowry conjugate acid

 

(b)(i)For the forward reaction from left to right, HCl donates a proton to form Cl and thus HCl is a proton donor .

It is a Bronsted-Lowry acid

 

  • For the backward reaction from right to left, Cl gains a proton to form HCl and thus Cl is an opposite proton acceptor.

 

It is a Bronsted-Lowry conjugate base.

 

-> 2M+(aq) + H2(g) -> M2+(aq) + H2(g) -> 2M3+(aq) + 3H2(g)

 

  1. Acids and bases show acidic and alkaline properties/characteristics only in water but not in other solvents e.g.

 

(a)Hydrogen chloride gas dissolves in water to form hydrochloric acid Hydrochloric acid dissociates/ionizes in water to free H+(aq)/H3O+(aq) ions. The free H3O+(aq) / H+(aq) ions are responsible for:

 

(i)turning blue litmus paper/solution red.

(ii)show pH value 1/2/3/4/5/6

(iii)are good electrolytes/conductors of electricity/undergo electrolysis.

 

(iv)react with metals to produce /evolve hydrogen gas and a salt. i.e.

Ionically:

-For a monovalent metal: 2M(s) + 2H+(aq)

-For a divalent metal:         M(s)   + 2H+(aq)

-For a trivalent metal:    2M(s)   + 6H+(aq)

 

Examples:
-For a monovalent metal: 2Na(s) + 2H+(aq) -> 2Na+(aq) + H2(g)
-For a divalent metal: Ca(s) + 2H+(aq) -> Ca2+(aq)  + H2(g)
-For a trivalent metal: 2Al(s) + 6H+(aq) -> 2Al3+(aq) + 3H2(g)

 

(v)react with metal carbonates and hhydrogen carbonates to produce /evolve carbon(IV)oxide gas ,water and a salt. i.e.

 

Ionically:

-For a monovalent metal: M2CO3(s)+ 2H+(aq) -> 2M+(aq) + H2O (l)+ CO2(g)

 

 

-For a divalent metal: MCO3(s)+ 2H+(aq) -> M2+(aq) + H2O (l)+ CO2(g)

 

 

Examples:

-For a monovalent metal: K2CO3(s)+ 2H+(aq) -> 2K+(aq) + H2O (l)+ CO2(g)

 

 

-For a divalent metal: ZnCO3(s)+ 2H+(aq) -> Zn2+(aq) + H2O (l)+ CO2(g)

 

 

 

(vi)neutralize metal oxides/hydroxides to salt and water only. i.e.

 

Ionically:

-For a monovalent metal: M2O(s) + 2H+(aq) -> 2M+(aq) + H2O (l)

 

-For a divalent metal: MO(s) + 2H+(aq) -> M2+(aq) + H2O (l)
M(OH) 2(s) + 2H+(aq) -> M2+(aq) + 2H2O(l)
-For a trivalent metal: M2O3(s) + 6H+(aq) -> 2M3+(aq) + 3H2O (l)
M(OH) 3(s) + 3H+(aq) -> M3+(aq) + 3H2O(l)

 

Examples:

-For a monovalent metal: K2O(s) + 2H+(aq) -> 2K+(aq) + H2O (l)

 

 

-For a divalent metal: ZnO (s) + 2H+(aq) -> Zn2+(aq) + H2O (l)

 

 

 

(b)Hydrogen chloride gas dissolves in methylbenzene /benzene but does not dissociate /ionize into free ions.

It exists in molecular state showing none of the above properties.

 

(c)Ammonia gas dissolves in water to form aqueous ammonia which dissociate/ionize to free NH4+ (aq) and OH(aq) ions.

This dissociation/ionization makes aqueous ammonia to:

 

(i)turn litmus paper/solution blue.

 

(ii)have pH 8/9/10/11

 

(iii)be a good electrical conductor

 

(iv)react with acids to form ammonium salt and water only.

 

NH4OH(aq) +        HCl(aq) ->   NH4Cl(aq)   + H2O(l)

 

(d)Ammonia gas dissolves in methylbenzene/benzene /kerosene but does not dissociate into free ions therefore existing as molecules

 

  1. Solvents are either polar or non-polar.

 

A polar solvent is one which dissolves ionic compounds and other polar solvents. Water is polar solvent that dissolves ionic and polar substance by surrounding the free ions as below:

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

H ð+

 

H ð+

 

O

 

ð-

 

 

 

H ð+

 

 

H ð+

 

 

 

H ð+

 

 

O

 

ð-

 

 

H ð+

 

 

H ð+

 

 

H ð+

 

 

 

 

 

 

 

O ð- H+   Oð- O ð- Cl O ð-
H ð+ Hð+
H ð+ O ð- H ð+ H+ H ð+
H ð+ H ð+ O ð-

 

 

 

 

Cl

 

 

 

Beaker

 

 

Cl        H+       water

H+

Cl        H+ Free ions

 

 

 

Note:Water is polar .It is made up of :

 

Oxygen atom is partially negative and two hydrogen atoms which are partially positive.

They surround the free H+ and Cl ions.

 

A non polar solvent is one which dissolved non-polar substances and covalent compounds.

 

If a polar ionic compound is dissolved in non-polar solvent ,it does not ionize/dissociate into free ions as below:

 

H-Cl

 

H-Cl

methyl benzene

H-Cl

H-Cl

 

 

 

Covalent bond

 

  1. Some acids and bases are strong while others are weak.

 

(a)A strong acid/base is one which is fully/wholly/completely dissociated / ionized into many free H+ /OH ions i.e.

  1. Strong acids exists more as free H+ ions than molecules. e.g.

 

HCl(aq) H+(aq) + Cl (aq)
(molecules) (cation) (anion)
HNO3(aq) H+(aq) + NO3 (aq)
(molecules) (cation) (anion)
H2SO4(aq) 2H+(aq) + SO4 2-(aq)
(molecules) (cation) (anion)
  1. Strong bases/alkalis exists more as free OH ions than molecules. e.g.

 

KOH(aq) K+(aq) + OH (aq)
(molecules) (cation) (anion)
NaOH(aq) Na+(aq) + OH(aq)
(molecules) (cation) (anion)

 

  • A weak base/acid is one which is partially /partly dissociated /ionized in water into free OH (aq) and H+(aq) ions.

 

  1. Weak acids exists more as molecules than as free H+ ions. e.g.

 

CH3COOH(aq) H+(aq) + CH3COO (aq)
(molecules) (cation) (anion)
H3PO4(aq) 3H+(aq) + PO4 3-(aq)
(molecules) (cation) (anion)
H2CO3(aq) 2H+(aq) + CO3 2-(aq)
(molecules) (cation) (anion)

 

  1. Weak bases/alkalis exists more as molecules than free OH ions. e.g.

 

NH4OH(aq) NH4 +(aq)+ OH (aq)
(molecules) (cation) (anion)
Ca(OH)2(aq) Ca2+(aq) + 2OH(aq)
(molecules) (cation) (anion)
Mg(OH)2(aq) Mg2+(aq) + 2OH(aq)
(molecules) (cation) (anion)

 

 

  1. The concentration of an acid/base/alkali is based on the number of moles of acid/bases dissolved in a decimeter(litre)of the solution.

 

An acid/base/alkali with more acid/base/alkali in a decimeter(litre) of solution is said to be concentrated while that with less is said to be dilute.

 

  1. (a) (i)strong acids have pH 1/2/3 while weak acids have high pH 4/5/6.

 

(ii)a neutral solution have pH 7

 

(iii)strong alkalis/bases have pH 12/13/14 while weak bases/alkalis have pH 11/10 /9 / 8.

 

  • pH is a measure of H+(aq) concentration in a solution. The higher the H+(aq)ions concentration ;

 

-the higher the acidity

 

-the lower the pH

 

-the lower the concentration of OH(aq) -the lower the alkalinity

 

At pH 7 , a solution has equal concentration of H+(aq) and OH(aq). Beyond pH 7,the concentration of the OH(aq) increases as the H+(aq) ions decreases.

 

10.(a) When acids /bases dissolve in water, the ions present in the solution conduct electricity.

 

The more the dissociation the higher the yield of ions and the greater the electrical conductivity of the solution.

 

A compound that conducts electricity in an electrolyte and thus a compound showing high electrical conductivity is a strong electrolyte while a compound showing low electrical conductivity is a weak electrolyte.

 

  • Practically, a bright light on a bulb ,a high voltage reading from a voltmeter high ammeter reading from an ammeter, a big deflection on a galvanometer is an indicator of strong electrolyte(acid/base) and the opposite for weak electrolytes(acids/base)

 

  1. Some compounds exhibit/show both properties of acids and bases/alkalis.

A substance that reacts with both acids and bases is said to be amphotellic.

The examples below show the amphotellic properties of:

 

  • Zinc (II)oxide(ZnO) and Zinc hydroxide(Zn(OH)2)

 

(i)When ½ spatula full of Zinc(II)oxide is placed in a boiling tube containing 10cm3 of either 2M nitric(V)acid or 2M sodium hydroxide hydroxide solution, it dissolves on both the acid and the alkali/base to form a colourless solution. i.e.

 

 

  • when reacting with nitric(V)acid, the oxide shows basic properties by reacting with an acid to form a simple salt and water

 

Basic oxide +   Acid -> salt   + water
Examples:
Chemical equation
ZnO(s)   +   2HNO3(aq) -> Zn(NO3) 2 (aq) + H2O(l)
ZnO(s) + 2HCl(aq) -> ZnCl2 (aq) + H2O(l)
ZnO(s)   +   H2SO4(aq) -> ZnSO4 (aq) + H2O(l)
Ionic equation
ZnO(s) + 2H+ (aq) -> Zn 2+ (aq) + H2O(l)

 

  • when reacting with sodium hydroxide, the oxide shows acidic properties by reacting with a base to form a complex

 

Basic oxide         +     Base/alkali   + Water        ->      Complex salt

Examples:

Chemical equation

 

1.When Zinc oxide is reacted with sodium hydroxide the complex salt is sodium tetrahydroxozincate(II) complex salt.

ZnO(s)   +    2NaOH(aq) +    H2O(l)  ->     Na2Zn(OH) 4(aq)

 

2.When Zinc oxide is reacted with potassium hydroxide the complex salt is potassium tetrahydroxozincate(II) complex salt.

ZnO(s)   +    2KOH(aq) +    H2O(l)  ->     K2Zn(OH) 4(aq)

 

Ionic equation

ZnO(s)   +    2OH(aq)   +    H2O(l)  ->     2[Zn(OH) 4]2- (aq)

 

(ii)When Zinc(II)hydroxide is placed in a boiling tube containing 10cm3 of either 2M nitric(V)acid or 2M sodium hydroxide hydroxide solution, it dissolves on both the acid and the alkali/base to form a colourless solution. i.e.

 

  • when reacting with nitric(V)acid, the hydroxide shows basic It reacts with an acid to form a simple salt and water only.

 

Basic hydroxide +   Acid -> salt   +water
Examples:
Chemical equation
Zn(OH) 2 (s)   + 2HNO3(aq) -> Zn(NO3) 2 (aq)  +   2H2O(l)

 

 

Zn(OH) 2 (s) + 2HCl(aq) -> ZnCl2 (aq) + 2H2O(l)
Zn(OH) 2 (s) + H2SO4(aq) -> ZnSO4 (aq) + 2H2O(l)
Ionic equation
Zn(OH) 2 (s) + 2H+ (aq) -> Zn 2+ (aq) +   2H2O(l)

 

  • when reacting with sodium hydroxide, the hydroxide shows acidic properties by reacting with a base to form a complex

 

Basic hydroxide        +        Base/alkali     ->      Complex salt

Examples:

Chemical equation

 

1.When Zinc hydroxide is reacted with sodium hydroxide the complex salt is sodium tetrahydroxozincate(II) complex salt.

 

Zn(OH) 2 (s)   +    2NaOH(aq) ->      Na2Zn(OH) 4(aq)

 

2.When Zinc hydroxide is reacted with potassium hydroxide the complex salt is potassium tetrahydroxozincate(II) complex salt.

 

Zn(OH) 2 (s) + 2KOH(aq) -> K2Zn(OH) 4(aq)
Ionic equation
Zn(OH) 2 (s) + 2OH(aq) -> 2[Zn(OH) 4]2- (aq)

 

(b)  Lead (II)oxide(PbO) and Lead(II) hydroxide (Pb(OH)2)

 

(i)When ½ spatula full of Lead(II)oxide is placed in a boiling tube containing 10cm3 of either 2M nitric(V)acid or 2M sodium hydroxide hydroxide solution, it dissolves on both the acid and the alkali/base to form a colourless solution. i.e.

 

  • when reacting with nitric(V)acid, the oxide shows basic properties by reacting with an acid to form a simple salt and water All other Lead salts are insoluble.

 

Chemical equation
PbO(s)   +   2HNO3(aq) -> Pb(NO3) 2 (aq) +   H2O(l)
Ionic equation
PbO(s)   +   2H+ (aq) -> Pb 2+ (aq)+ H2O(l)

 

  • when reacting with sodium hydroxide, the oxide shows acidic properties by reacting with a base to form a complex

 

Chemical equation

 

1.When Lead(II) oxide is reacted with sodium hydroxide the complex salt is sodium tetrahydroxoplumbate(II) complex salt.

PbO(s)    +    2NaOH(aq) +    H2O(l)   ->     Na2Pb(OH) 4(aq)

 

2.When Lead(II) oxide is reacted with potassium hydroxide the complex salt is potassium tetrahydroxoplumbate(II) complex salt.

PbO(s)    +    2KOH(aq) +    H2O(l)  ->    K2Pb(OH) 4(aq)

 

Ionic equation

PbO(s)    +    2OH(aq)   +    H2O(l)  ->    2[Pb(OH) 4]2- (aq)

 

(ii)When Lead(II)hydroxide is placed in a boiling tube containing 10cm3 of either 2M nitric(V)acid or 2M sodium hydroxide hydroxide solution, it dissolves on both the acid and the alkali/base to form a colourless solution. i.e.

 

  • when reacting with nitric(V)acid, the hydroxide shows basic It reacts with the acid to form a simple salt and water only.

 

Chemical equation
Pb(OH) 2 (s) + 2HNO3(aq) -> Pb(NO3) 2 (aq) +   2H2O(l)
Ionic equation
Pb(OH) 2 (s) + 2H+ (aq) -> Pb 2+ (aq)+ 2H2O(l)

 

  • when reacting with sodium hydroxide, the hydroxide shows acidic It reacts with a base to form a complex salt.

 

Chemical equation

 

1.When Lead(II) hydroxide is reacted with sodium hydroxide the complex salt is sodium tetrahydroxoplumbate(II) complex salt.

 

Pb(OH) 2 (s)   +    2NaOH(aq) ->      Na2Pb(OH) 4(aq)

 

2.When Lead(II) hydroxide is reacted with potassium hydroxide the complex salt is potassium tetrahydroxoplumbate(II) complex salt.

 

Pb(OH) 2 (s) + 2KOH(aq) -> K2Pb(OH) 4(aq)
Ionic equation
Pb(OH) 2 (s) + 2OH(aq) -> 2[Pb(OH) 4]2- (aq)

 

 

(c)Aluminium(III)oxide(Al2O3) and Aluminium(III)hydroxide(Al(OH)3)

 

(i)When ½ spatula full of Aluminium(III)oxide is placed in a boiling tube containing 10cm3 of either 2M nitric(V)acid or 2M sodium hydroxide hydroxide solution, it dissolves on both the acid and the alkali/base to form a colourless solution. i.e.

 

  • when reacting with nitric(V)acid, the oxide shows basic properties by reacting with an acid to form a simple salt and water

 

Chemical equation
Al2O3 (s) + 6HNO3(aq) -> Al(NO3)3 (aq) + 3H2O(l)
Al2O3 (s) + 6HCl(aq) -> AlCl3 (aq) + 3H2O(l)
Al2O3 (s) + 3H2SO4(aq) -> Al2(SO4)3 (aq) + 3H2O(l)
Ionic equation
Al2O3 (s) + 3H+ (aq) -> Al 3+ (aq)+ 3H2O(l)

 

  • when reacting with sodium hydroxide, the oxide shows acidic properties by reacting with a base to form a complex

 

Chemical equation

 

1.When Aluminium(III) oxide is reacted with sodium hydroxide the complex salt is sodium tetrahydroxoaluminate(III) complex salt.

 

Al2O3 (s)   +    2NaOH(aq) +      3H2O(l)   ->     2NaAl(OH) 4(aq)

 

2.When Aluminium(III) oxide is reacted with potassium hydroxide the complex salt is potassium tetrahydroxoaluminate(II) complex salt.

 

Al2O3 (s) + 2KOH(aq) + 3H2O(l) -> 2NaAl(OH) 4(aq)
Ionic equation
Al2O3 (s) + 2OH(aq)  + 3H2O(l) -> 2[Al(OH) 4] (aq)

 

(ii)When Aluminium(III)hydroxide is placed in a boiling tube containing 10cm3 of either 2M nitric(V)acid or 2M sodium hydroxide hydroxide solution, it dissolves on both the acid and the alkali/base to form a colourless solution. i.e.

 

  • when reacting with nitric(V)acid, the hydroxide shows basic It reacts with the acid to form a simple salt and water only.

 

Chemical equation

Al(OH) 3 (s)    +    3HNO3(aq)   ->    Al(NO3)3 (aq)        +                                    3H2O(l)

 

Al(OH)3 (s) + 3HCl(aq) -> AlCl3 (aq) + 3H2O(l)
2Al(OH)3 (s) + 3H2SO4(aq) -> Al2(SO4)3 (aq) + 3H2O(l)
Ionic equation
Al(OH)3 (s) + 3H+ (aq) -> Al 3+ (aq)+ 3H2O(l)

 

  • when reacting with sodium hydroxide, the hydroxide shows acidic It reacts with a base to form a complex salt.

 

Chemical equation

 

1.When aluminium(III) hydroxide is reacted with sodium hydroxide the complex salt is sodium tetrahydroxoaluminate(III) complex salt.

Al(OH) 3 (s)   +    NaOH(aq) ->      NaAl(OH) 4(aq)

 

2.When aluminium(III) hydroxide is reacted with potassium hydroxide the complex salt is potassium tetrahydroxoaluminate(III) complex salt.

Al(OH) 3 (s) + KOH(aq) -> KAl(OH) 4(aq)
Ionic equation
Al(OH) 3 (s) + OH(aq) -> [Al(OH) 4] (aq)
Summary of amphotellic oxides/hydroxides
Oxide Hydroxide Formula of simple salt Formula of complex salt
from nitric (V)acid from sodium hydroxide
ZnO Zn(OH)2 Zn(NO3)2 Na2Zn(OH)4
[Zn(OH)4]2-(aq)
Sodium tetrahydroxozincate(II)
PbO Pb(OH)2 Pb(NO3)2 Na2Pb(OH)4
[Pb(OH)4]2-(aq)
Sodium tetrahydroxoplumbate(II)
Al2O3 Al(OH) 3 Al(NO3)3 NaAl(OH)4
[Al(OH)4](aq)
Sodium tetrahydroxoaluminate(II)

 

12.(a) A salt is an ionic compound formed when the cation from a base combine with the anion derived from an acid.

 

 

A salt is therefore formed when the hydrogen ions in an acid are replaced wholly/fully or partially/partly ,directly or indirectly by a metal or ammonium radical.

 

  • The number of ionizable/replaceable hydrogen in an acid is called basicity of an acid.

Some acids are therefore:

(i)monobasic acids generally denoted HX e.g.

HCl, HNO3,HCOOH,CH3COOH.

 

(ii)dibasic acids ; generally denoted H2X e.g.

H2SO4, H2SO3, H2CO3,HOOCOOH.

 

(iii)tribasic acids ; generally denoted H3X e.g.

H3PO4.

 

(c) Some salts are normal salts while other are acid salts.

 

(i)A normal salt is formed when all the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical.

 

(ii)An acid salt is formed when part/portion the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical.

 

Table showing normal and acid salts derived from common acids

 

Acid  name Chemical Basicity Normal salt Acid salt
formula
Hydrochloric acid HCl Monobasic Chloride(Cl) None
Nitric(V)acid HNO3 Monobasic Nitrate(V)(NO3 ) None
Nitric(III)acid HNO2 Monobasic Nitrate(III)(NO2 ) None
Sulphuric(VI)acid H2SO4 Dibasic Sulphate(VI) (SO4 2-) Hydrogen sulphate(VI)
(HSO4 )
Sulphuric(IV)acid H2SO3 Dibasic Sulphate(IV) (SO3 2-) Hydrogen sulphate(IV)
(HSO3 )
Carbonic(IV)acid H2CO3 Dibasic Carbonate(IV)(CO3 2-) Hydrogen carbonate(IV)
(HCO3 )
Phosphoric(V) H3PO4 Tribasic Phosphate(V)(PO4 3-) Dihydrogen

 

 

acid phosphate(V)
(H2PO42-)
Hydrogen
diphosphate(V)
(HP2O42-)

 

The table below show shows some examples of salts.

Base/alkali Cation Acid Anion Salt Chemical name of salts
NaOH + HCl NaCl Sodium(I)chloride
Na Cl
Mg(OH)2 Mg2+ H2SO4 SO 2 MgSO4 Magnesium sulphate(VI)
4 Mg(HSO4)2 Magnesium hydrogen
sulphate(VI)
Pb(OH)2 Pb2+ HNO3 NO Pb(NO3)2 Lead(II)nitrate(V)
3
Ba(OH)2 Ba2+ HNO3 NO Ba(NO3)2 Barium(II)nitrate(V)
3
Ca(OH)2 Ba2+ H2SO4 SO 2 MgSO4 Calcium sulphate(VI)
4
NH4OH NH + H3PO4 PO 3 (NH4 )3PO4 Ammonium phosphate(V)
4 4 (NH4 )2HPO4 Diammonium phosphate(V)
Ammonium diphosphate(V)
NH4 H2PO4
KOH K+ H3PO4 PO 3 K3PO4 Potassium phosphate(V)
4
Al(OH)3 Al3+ H2SO4 SO 2 Al2(SO4)2 Aluminium(III)sulphate(VI)
4
Fe(OH)2 Fe2+ H2SO4 SO 2 FeSO4 Iron(II)sulphate(VI)
4
Fe(OH)3 Fe3+ H2SO4 SO 2 Fe2(SO4)2 Iron(III)sulphate(VI)
4

 

  • Some salts undergo hygroscopy, deliquescence and efflorescence.

 

 

  • Hygroscopic salts /compounds are those that absorb water from the atmosphere but do not form a solution.

 

Some salts which are hygroscopic include anhydrous copper(II)sulphate(VI), anhydrous cobalt(II)chloride, potassium nitrate(V) common table salt.

 

(ii)Deliquescent salts /compounds are those that absorb water from the atmosphere and form a solution.

 

Some salts which are deliquescent include: Sodium nitrate(V),Calcium chloride, Sodium hydroxide, Iron(II)chloride, Magnesium chloride.

 

(iii)Efflorescent salts/compounds are those that lose their water of crystallization to the atmosphere.

 

Some salts which effloresces include: sodium carbonate decahydrate, Iron(II)sulphate(VI)heptahydrate, sodium sulphate (VI)decahydrate.

 

(e)Some salts contain water of crystallization.They are hydrated.Others do not contain water of crystallization. They are anhydrous.

 

Table showing some hydrated salts.

Name of hydrated salt Chemical formula
Copper(II)sulphate(VI)pentahydrate CuSO4.5H2O
Aluminium(III)sulphate(VI)hexahydrate Al2 (SO4) 3.6H2O
Zinc(II)sulphate(VI)heptahydrate ZnSO4.7H2O
Iron(II)sulphate(VI)heptahydrate FeSO4.7H2O
Calcium(II)sulphate(VI)heptahydrate CaSO4.7H2O
Magnesium(II)sulphate(VI)heptahydrate MgSO4.7H2O
Sodium sulphate(VI)decahydrate Na2SO4.10H2O
Sodium carbonate(IV)decahydrate Na2CO3.10H2O
Potassium carbonate(IV)decahydrate K2CO3.10H2O
Potassium sulphate(VI)decahydrate K2SO4.10H2O

 

(f)Some salts exist as a simple salt while some as complex salts. Below are some complex salts.

 

Table of some complex salts

 

Name of complex salt Chemical formula Colour of the complex salt
Tetraamminecopper(II)sulphate(VI) Cu(NH3) 4 SO4 H2O Royal/deep blue solution
Tetraamminezinc(II)nitrate(V) Zn(NH3) 4 (NO3 )2 Colourless solution
Tetraamminecopper(II) nitrate(V) Cu(NH3) 4 (NO3 )2 Royal/deep blue solution
Tetraamminezinc(II)sulphate(VI) Zn(NH3) 4 SO4 Colourless solution

 

(g)Some salts exist as two salts in one. They are called double salts.

 

Table of some double salts

Name of double salts Chemical formula
Trona(sodium sesquicarbonate) Na2CO3 NaHCO3.2H2O
Ammonium iron(II)sulphate(VI) FeSO4(NH4) 2SO4.2H2O
Ammonium aluminium(III)sulphate(VI) Al2(SO4) 3(NH4) 2SO4.H2O

 

(h)Some salts dissolve in water to form a solution. They are said to be soluble.

Others do not dissolve in water. They form a suspension/precipitate in water.

 

Table of solubility of salts

 

Soluble salts Insoluble salts
All nitrate(V)salts
All sulphate(VI)/SO42- salts except Barium(II) sulphate(VI)/BaSO4
Calcium(II) sulphate(VI)/CaSO4
Lead(II) sulphate(VI)/PbSO4
All sulphate(IV)/SO32- salts except Barium(II) sulphate(IV)/BaSO3
Calcium(II) sulphate(IV)/CaSO3
Lead(II) sulphate(IV)/PbSO3
All chlorides/Cl except Silver chloride/AgCl
Lead(II)chloride/PbCl2(dissolves in hot
water)
All phosphate(V)/PO4 3-

All sodium,potassium and ammonium

 

salts

All hydrogen carbonates/HCO3

All hydrogen sulphate(VI)/ HSO4

Sodium carbonate/Na2CO3, except All carbonates
potassium carbonate/ K2CO3,
ammonium carbonate (NH4) 2CO3
All alkalis(KOH,NaOH, NH4OH) except All bases

 

13 Salts can be prepared in a school laboratory by a method that uses its solubility in water.

(a) Soluble salts may be prepared by using any of the following methods:

 

(i)Direct displacement/reaction of a metal with an acid.

 

 

By reacting a metal higher in the reactivity series than hydrogen with a dilute acid,a salt is formed and hydrogen gas is evolved.

Excess of the metal must be used to ensure all the acid has reacted.

When effervescence/bubbling /fizzing has stopped ,excess metal is filtered.

 

The filtrate is heated to concentrate then allowed to crystallize.

 

Washing with distilled water then drying between filter papers produces a sample crystal of the salt. i.e.

M(s) + H2X-> MX(aq)  +   H2(g)
Examples
Mg(s) + H2SO4(aq) -> MgSO4 (aq) + H2(g)
Zn(s) + H2SO4(aq) -> ZnSO4 (aq) + H2(g)
Pb(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2(g)
Ca(s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2(g)
Mg(s) + 2HNO3(aq) -> Mg(NO3) 2(aq) + H2(g)
Mg(s) + 2HCl(aq) -> MgCl 2(aq) + H2(g)
Zn(s) + 2HCl(aq) -> ZnCl 2(aq) + H2(g)

 

(ii)Reaction of an insoluble base with an acid

 

By adding an insoluble base (oxide/hydroxide )to a dilute acid until no more dissolves, in the acid,a salt and water are formed. Excess of the base is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter papers e.g.

PbO(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l)
Pb(OH)2(s)  + 2HNO3(aq) -> Pb(NO3) 2(aq)   + 2H2O (l)
CaO (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l)
MgO (s) + 2HNO3(aq) -> Mg(NO3) 2(aq) + H2O (l)
MgO (s) + 2HCl(aq) -> MgCl 2(aq) + H2O (l)
ZnO (s) + 2HCl(aq) -> ZnCl 2(aq) + H2O (l)
Zn(OH)2(s)  + 2HNO3(aq) -> Zn(NO3) 2(aq)  + 2H2O (l)
CuO (s) + 2HCl(aq) -> CuCl 2(aq) + H2O (l)
CuO (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l)
Ag2O(s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l)
Na2O(s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l)

 

(iii)reaction of insoluble /soluble carbonate /hydrogen carbonate with an acid. By adding an excess of a soluble /insoluble carbonate or hydrogen carbonate to adilute acid, effervescence /fizzing/bubbling out of carbon(IV)oxide gas shows the reaction is taking place. When effervescence /fizzing/bubbling out of the gas is over, excess of the insoluble carbonate is filtered off. The filtrate is heated to

 

 

concentrate ,allowed to crystallize then washed with distilled water before drying between filter paper papers e.g.

PbCO3 (s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l)+ CO2(g)
ZnCO3 (s) + 2HNO3(aq) -> Zn(NO3) 2(aq) + H2O (l)+ CO2(g)
CaCO3 (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l)+ CO2(g)
MgCO3 (s) + H2SO4(aq) -> MgSO4(aq) + H2O (l)+ CO2(g)
Cu CO3 (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l) + CO2(g)
Ag2CO3 (s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l) + CO2(g)
Na2CO3 (s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l) + CO2(g)
K2CO3 (s) + 2HCl(aq) -> 2KCl(aq) + H2O (l) + CO2(g)
NaHCO3 (s) + HNO3(aq) -> NaNO3(aq) + H2O (l) + CO2(g)
KHCO3 (s) + HCl(aq) -> KCl(aq) + H2O (l) + CO2(g)

(iv)neutralization/reaction of soluble base/alkali with dilute acid

 

By adding an acid to a burette into a known volume of an alkali with 2-3 drops of an indicator, the colour of the indicator changes when the acid has completely reacted with an alkali at the end point. The procedure is then repeated without the indicator .The solution mixture is then heated to concentrate , allowed to crystallize ,washed with distilled water before drying with filter papers. e.g.

 

NaOH (aq) + HNO3(aq) -> NaNO3(aq) + H2O (l)
KOH (aq) + HNO3(aq) -> KNO3(aq) + H2O (l)
KOH (aq) + HCl(aq) -> KCl(aq) + H2O (l)
2KOH (aq) + H2SO4(aq) -> K2SO4(aq) + 2H2O (l)
2 NH4OH (aq) + H2SO4(aq) -> (NH4)2SO4(aq) + 2H2O (l)
NH4OH (aq) + HNO3(aq) -> NH4NO3(aq) + H2O (l)

 

(iv)direct synthesis/combination.

 

When a metal burn in a gas jar containing a non metal , the two directly combine to form a salt. e.g.

2Na(s) + Cl2(g) -> 2NaCl(s)
2K(s) + Cl2(g) -> 2KCl(s)
Mg(s) + Cl2(g) -> Mg Cl2 (s)
Ca(s) + Cl2(g) -> Ca Cl2 (s)

 

Some salts once formed undergo sublimation and hydrolysis. Care should be taken to avoid water/moisture into the reaction flask during their preparation.Such salts include aluminium(III)chloride(AlCl3) and iron (III)chloride(FeCl3)

 

  1. Heated aluminium foil reacts with chlorine to form aluminium(III)chloride that sublimes away from the source of heating then deposited as solid again

 

2Al(s)                +          3Cl2(g)            ->          2AlCl3 (s/g)

 

Once formed aluminium(III)chloride hydrolyses/reacts with water vapour / moisture present to form aluminium hydroxide solution and highly acidic fumes of hydrogen chloride gas.

AlCl3(s)+       3H2 O(g)        ->          Al(OH)3 (aq) + 3HCl(g)

 

  1. Heated iron filings reacts with chlorine to form iron(III)chloride that sublimes away from the source of heating then deposited as solid again

2Fe(s)                +          3Cl2(g)            ->          2FeCl3 (s/g)

 

Once formed , aluminium(III)chloride hydrolyses/reacts with water vapour / moisture present to form aluminium hydroxide solution and highly acidic fumes of hydrogen chloride gas.

FeCl3(s)+       3H2 O(g)        ->          Fe(OH)3 (aq) + 3HCl(g)

 

(b)Insoluble salts can be prepared by reacting two suitable soluble salts to form one soluble and one insoluble. This is called double decomposition or precipitation. The mixture is filtered and the residue is washed with distilled water then dried.

CuSO4(aq) + Na2CO3 (aq) -> CuCO3 (s) + Na2 SO4(aq)
BaCl2(aq) + K2SO4 (aq) -> BaSO4 (s) + 2KCl (aq)
Pb(NO3)2(aq) + K2SO4 (aq) -> PbSO4 (s) + 2KNO3 (aq)
2AgNO3(aq) + MgCl2 (aq) -> 2AgCl(s) + Mg(NO3)2 (aq)
Pb(NO3)2(aq) + (NH4) 2SO4 (aq) -> PbSO4 (s) + 2NH4NO 3(aq)
BaCl2(aq) + K2SO3 (aq) -> BaSO3 (s) + 2KCl (aq)

 

  1. Salts may lose their water of crystallization , decompose ,melt or sublime on heating on a Bunsen burner flame.

 

The following shows the behavior of some salts on heating gently /or strongly in a laboratory school burner:

 

(a)effect of heat on chlorides

 

All chlorides have very high melting and boiling points and therefore are not affected by laboratory heating except ammonium chloride. Ammonium chloride sublimes on gentle heating. It dissociate into the constituent ammonia and hydrogen chloride gases on strong heating.

NH4Cl(s)   NH4Cl(g)   NH3(g) + HCl(g)

(sublimation)                     (dissociation)

 

(b)effect of heat on nitrate(V)

 

  • Potassium nitrate(V)/KNO3 and sodium nitrate(V)/NaNO3 decompose on heating to form Potassium nitrate(III)/KNO2 and sodium nitrate(III)/NaNO2 and

 

producing Oxygen gas in each case.

2KNO3 (s) -> 2KNO2(s) + O2(g) 2NaNO3 (s) -> 2NaNO2(s) + O2(g)

 

(ii)Heavy metal nitrates(V) salts decompose on heating to form the oxide and a mixture of brown acidic nitrogen(IV)oxide and oxygen gases. e.g.

2Ca(NO3)2 (s)             ->    2CaO(s)    +   4NO2(g)  +      O2(g)

2Mg(NO3)2(s)            ->    2MgO(s)    +   4NO2(g)   +      O2(g)

2Zn(NO3)2(s)              ->    2ZnO(s)    +   4NO2(g)  +      O2(g)

2Pb(NO3)2(s)               ->    2PbO(s)    +   4NO2(g)  +      O2(g)

2Cu(NO3)2(s)              ->    2CuO(s)    +   4NO2(g)   +      O2(g)

2Fe(NO3)2(s)               ->    2FeO(s)    +   4NO2(g)  +      O2(g)

 

(iii)Silver(I)nitrate(V) and mercury(II) nitrate(V) are lowest in the reactivity series. They decompose on heating to form the metal(silver and mercury)and the Nitrogen(IV)oxide and oxygen gas. i.e.

 

 

2AgNO3(s)    ->    2Ag (s)    +   2NO2(g)   +    O2(g)

2Hg(NO3)2 (s)             ->    2Hg (s)    +   4NO2(g)   +

 

 

O2(g)

 

 

 

(iv)Ammonium nitrate(V) and Ammonium nitrate(III) decompose on heating to Nitrogen(I)oxide(relights/rekindles glowing splint) and nitrogen gas respectively.Water is also formed.i.e.

 

 

NH4NO3(s)          ->         N2O (g)

NH4NO2(s)          ->         N2 (g)

 

+     H2O(l)

+       H2O(l)

 

 

 

(c) effect of heat on nitrate(V)

 

Only Iron(II)sulphate(VI), Iron(III)sulphate(VI) and copper(II)sulphate(VI) decompose on heating. They form the oxide, and produce highly acidic fumes of acidic sulphur(IV)oxide gas.

2FeSO4 (s) -> Fe2O3(s) + SO3(g) +   SO2(g)
Fe2(SO4) 3(s) -> Fe2O3(s)   +   SO3(g)
CuSO4 (s) -> CuO(s) + SO3(g)

 

  • effect of heat on carbonates(IV) and hydrogen carbonate(IV). (i)Sodium carbonate(IV)and potassium carbonate(IV)do not decompose on heating.

 

(ii)Heavy metal nitrate(IV)salts decompose on heating to form the oxide and produce carbon(IV)oxide gas. Carbon (IV)oxide gas forms a white precipitate

 

when bubbled in lime water. The white precipitate dissolves if the gas is in excess.

e.g. CuCO3 (s) -> CuO(s) + CO2(g)
CaCO3 (s) -> CaO(s) + CO2(g)
PbCO3 (s) -> PbO(s) + CO2(g)
FeCO3 (s) -> FeO(s) + CO2(g)
ZnCO3 (s) -> ZnO(s) + CO2(g)

 

(iii)Sodium hydrogen carbonate(IV) and Potassium hydrogen carbonate(IV)decompose on heating to give the corresponding carbonate (IV) and form water and carbon(IV)oxide gas. i.e.

 

2NaHCO 3(s)                   ->         Na2CO3(s)     +     CO2(g)   + H2O(l)

2KHCO 3(s)                        ->         K2CO3(s)         +     CO2(g)   + H2O(l)

 

  • Calcium hydrogen carbonate (IV) and Magnesium hydrogen carbonate(IV) decompose on heating to give the corresponding carbonate (IV) and form water

and carbon(IV)oxide gas. i. e.

Ca(HCO3) 2(aq)->CaCO3(s)+   CO2(g)  + H2O(l)

Mg(HCO3) 2(aq)                 ->         MgCO3(s)    +        CO2(g)   + H2O(l)

 

  1. Salts contain cation(positively charged ion) and anions(negatively charged ion).When dissolved in polar solvents/water.

 

The cation and anion in a salt is determined/known usually by precipitation of the salt using a precipitating reagent.

 

The colour of the precipitate is a basis of qualitative analysis of a compound.

 

16.Qualitative analysis is the process of identifying an unknown compound /salt by identifying the unique qualities of the salt/compound. It involves some of the following processes.

 

(a)Reaction of cation with sodium/potassium hydroxide solution.

 

Both sodium/potassium hydroxide solutions are precipitating reagents.

 

The alkalis produce unique colour of a precipitate/suspension when a few/three drops is added and then excess alkali is added to unknown salt/compound solution.

 

NB: Potassium hydroxide is not commonly used because it is more expensive than sodium hydroxide.

 

The table below shows the observations, inferences / deductions and explanations from the following test tube experiments:

 

Procedure

Put about 2cm3 of MgCl2, CaCl2, AlCl3, NaCl, KCl, FeSO4, Fe2(SO4) 3, CuSO4, ZnSO4NH4NO3, Pb(NO3) 2, Ba(NO3) 2 each into separate test tubes. Add three

 

drops of 2M sodium hydroxide solution then excess (2/3 the length of a standard test tube).

 

 

 

 

 

 

 

 

 

 

 

 

Observation Inference Explanation
No white Na+ and K+ Both Na+ and K+ ions react with OH from 2M
precipitate sodium hydroxide solution to form soluble
colourless solutions
Na+(aq) + OH(aq) -> NaOH(aq)
K+(aq) + OH(aq) -> KOH(aq)
No white NH4 + ions NH4 + ions react with 2M sodium hydroxide
precipitate then solution to produce pungent smelling ammonia
pungent smell of gas
ammonia /urine
NH4 + (aq) + OH(aq) -> NH3 (g) + H2O(l)
White precipitate Ba2+ ,Ca2+, Ba2+ ,Ca2+ and Mg2+ ions react with OH from
insoluble in Mg2+ ions 2M sodium hydroxide solution to form insoluble
excess white precipitate of their hydroxides.
Ba2+(aq) + 2OH(aq) -> Ba(OH) 2(s)
Ca2+(aq) + 2OH(aq) -> Ca(OH) 2(s)
Mg2+(aq) + 2OH(aq) -> Mg(OH) 2(s)
White precipitate Zn2+ ,Pb2+, Pb2+ ,Zn2+ and Al3+ ions react with OH from
soluble in excess Al3+ ions 2M sodium hydroxide solution to form insoluble
white precipitate of their hydroxides.
Zn2+(aq) + 2OH(aq) -> Zn(OH) 2(s)
Pb2+(aq) + 2OH(aq) -> Pb(OH) 2(s)
Al3+(aq) + 3OH(aq) -> Al(OH) 3(s)

 

 

The hydroxides formed react with more OH
ions to form complex salts/ions.
Zn(OH) 2(s) + 2OH(aq) -> [ Zn(OH) 4]2-(aq)
Pb(OH) 2(s) + 2OH(aq) -> [ Pb(OH) 4]2-(aq)
Al(OH) 3(s) + OH(aq) -> [ Al(OH) 4](aq)
Blue precipitate Cu2+ Cu2+ ions react with OH from 2M sodium
insoluble in hydroxide solution to form insoluble blue
excess precipitate of copper(II) hydroxide.
Cu2+(aq) + 2OH(aq) -> Cu(OH) 2(s)
Green Fe2+ Fe2+ ions react with OH from 2M sodium
precipitate hydroxide solution to form insoluble green
insoluble in precipitate of Iron(II) hydroxide.
excess
Fe2+ oxidized Fe2+(aq) + 2OH(aq) -> Fe(OH) 2(s)
On adding 3cm3 Hydrogen peroxide is an oxidizing agent that
of hydrogen to Fe3+ oxidizes green Fe2+ oxidized to brown Fe3+
peroxide,
brown/yellow Fe(OH) 2(s) + 2H+  -> Fe(OH) 3(aq)
solution formed
Brown Fe3+ Fe3+ ions react with OH from 2M sodium
precipitate hydroxide solution to form insoluble brown
insoluble in precipitate of Iron(II) hydroxide.
excess Fe3+(aq) + 3OH(aq) -> Fe(OH) 3(s)

 

 

 

 

(b)Reaction of cation with aqueous ammonia

 

 

 

Aqueous ammonia precipitating reagent that can be used to identify the cations present in a salt.

 

Like NaOH/KOH the OH ion in NH4OH react with the cation to form a characteristic hydroxide .

 

 

Below are the observations ,inferences and explanations of the reactions of aqueous ammonia with salts from the following test tube reactions.

 

Procedure

 

Put about 2cm3 of MgCl2, CaCl2, AlCl3, NaCl, KCl, FeSO4, Fe2(SO4) 3, CuSO4, ZnSO4NH4NO3, Pb(NO3) 2, Ba(NO3) 2 each into separate test tubes.

 

Add three drops of 2M aqueous ammonia then excess (2/3 the length of a standard test tube).

 

Observation Inference Explanation
No white Na+ and K+ NH4 +,Na+ and K+ ions react with OH from 2M
precipitate aqueous ammonia to form soluble colourless
solutions
NH4 + (aq) + OH(aq) -> NH4 +OH(aq)
Na+(aq) + OH(aq) -> NaOH(aq)
K+(aq) + OH(aq) -> KOH(aq)
White precipitate Ba2+ ,Ca2+, Ba2+ ,Ca2+,Mg2+ ,Pb2+ and Al3+, ions react with
insoluble in Mg2+ ,Pb2+, OH from 2M aqueous ammonia to form
excess Al3+, ions insoluble white precipitate of their hydroxides.
Pb2+ (aq) + 2OH(aq) -> Pb(OH) 2(s)
Al3+ (aq) + 3OH(aq) -> Al(OH) 3(s)
Ba2+ (aq) + 2OH(aq) -> Ba(OH) 2(s)
Ca2+ (aq) + 2OH(aq) -> Ca(OH) 2(s)
Mg2+ (aq) + 2OH(aq) -> Mg(OH) 2(s)
White precipitate Zn2+ ions Zn2+ ions react with OH from 2M aqueous
soluble in excess ammonia to form insoluble white precipitate of
Zinc hydroxide.
Zn2+(aq) + 2OH(aq) -> Zn(OH) 2(s)
The Zinc hydroxides formed react NH3(aq) to
form a complex salts/ions.
Zn(OH) 2(s) + 4NH3(aq)
->[ Zn(NH3) 4]2+(aq)+ 2OH(aq)
Blue precipitate Cu2+ Cu2+ ions react with OH from 2M aqueous

 

 

that dissolves in ammonia to form blue precipitate of copper(II)
excess ammonia hydroxide.
solution to form Cu2+(aq) + 2OH(aq) -> Cu(OH) 2(s)
a deep/royal
The copper(II) hydroxide formed react NH3(aq)
blue solution
to form a complex salts/ions.
Cu(OH) 2 (s) + 4NH3(aq)
->[ Cu(NH3) 4]2+(aq)+ 2OH(aq)
Green Fe2+ Fe2+ ions react with OH from 2M aqueous
precipitate ammonia to form insoluble green precipitate of
insoluble in Iron(II) hydroxide.
excess. Fe2+(aq) + 2OH(aq) -> Fe(OH) 2(s)
On adding 3cm3 Fe2+ oxidized Hydrogen peroxide is an oxidizing agent that
of hydrogen to Fe3+ oxidizes green Fe2+ oxidized to brown Fe3+
peroxide, Fe(OH) 2(s) + 2H+  -> Fe(OH) 3(aq)
brown/yellow
solution formed
Brown Fe3+ Fe3+ ions react with OH from 2M aqueous
precipitate ammonia to form insoluble brown precipitate of
insoluble in Iron(II) hydroxide.
excess Fe3+(aq) + 3OH(aq) -> Fe(OH) 3(s)

 

Note

 

  • Only Zn2+ ions/salts form a white precipitate that dissolve in excess of both 2M sodium hydroxide and 2M aqueous ammonia.

 

  • Pb2+ and Al3+ ions/salts form a white precipitate that dissolve in excess of 2M sodium hydroxide but not in 2M aqueous ammonia.

 

  • Cu2+ ions/salts form a blue precipitate that dissolve to form a deep/royal blue solution in excess of 2M aqueous ammonia but only blue insoluble precipitate in 2M sodium hydroxide

 

(c)Reaction of cation with Chloride (Cl)ions

 

All chlorides are soluble in water except Silver chloride and Lead (II)chloride (That dissolve in hot water).When a soluble chloride like NaCl, KCl, NH4Cl is

 

 

added to about 2cm3 of a salt containing Ag+ or Pb2+ions a white precipitate of AgCl or PbCl2 is formed. The following test tube reactions illustrate the above.

 

Experiment

 

Put about 2cm3 of silver nitrate(V) andLead(II)nitrate(V)solution into separate test tubes. Add five drops of NaCl /KCl / NH4Cl/HCl. Heat to boil.

 

 

 

 

Observation Inference Explanation
(i)White precipitate Ag+ ions Ag+ ions reacts with Cl ions from a
does not dissolve on soluble chloride salt to form a white
heating precipitate of AgCl
(ii)White precipitate Pb2+ ions Pb2+ ions reacts with Cl ions from a
dissolve on heating soluble chloride salt to form a white
precipitate of PbCl2. PbCl2 dissolves
on heating.
Note

 

Both Pb2+ and Al3+ ions forms an insoluble white precipitate in excess aqueous ammonia. A white precipitate on adding Cl ions/salts shows Pb2+. No white precipitate on adding Cl ions/salts shows Al3+.

 

Adding a chloride/ Cl ions/salts can thus be used to separate the identity of Al3+ and Pb2+.

 

(d)Reaction of cation with sulphate(VI)/SO42- and sulphate(IV)/SO32-  ions

 

All sulphate(VI) and sulphate(IV)/SO32- ions/salts are soluble/dissolve in water except Calcium sulphate(VI)/CaSO4, Calcium sulphate(IV)/CaSO3, Barium sulphate(VI)/BaSO4, Barium sulphate(IV)/BaSO3, Lead(II) sulphate(VI)/PbSO4

 

and Lead(II) sulphate(IV)/PbSO3.When a soluble sulphate(VI)/SO42- salt like Na2SO4, H2SO4, (NH4)2SO4 or Na2SO3 is added to a salt containing Ca2+, Pb2+,

Ba2+ ions, a white precipitate is formed.

The following test tube experiments illustrate the above.

 

Procedure

Place about 2cm3 of Ca(NO3)2, Ba(NO3)2, BaCl2 and Pb(NO3)2, in separate boiling tubes. Add six drops of sulphuric(VI)acid /sodium sulphate(VI)/ammonium sulphate(VI)solution. Repeat with six drops of sodium sulphate(IV).

 

Observation Inference Explanation
White Ca2+, Ba2+, CaSO3 and CaSO4 do not form a thick precipitate as

 

 

precipitate Pb2+ ions they are sparingly soluble.
Ca2+(aq)+ SO3 2-(aq) -> CaSO3(s)
Ca2+(aq)+ SO4 2-(aq) -> CaSO4(s)
Ba2+(aq)+ SO3 2-(aq) -> BaSO3(s)
Ba2+(aq)+ SO4 2-(aq) -> BaSO4(s)
Pb2+(aq)+ SO32-(aq) -> PbSO3(s)
Pb2+(aq)+ SO42-(aq) -> PbSO4(s)

(e)Reaction of cation with carbonate(IV)/CO32-  ions

 

All carbonate salts are insoluble except sodium/potassium carbonate(IV) and ammonium carbonate(IV).

 

They dissociate /ionize to release CO32- ions. CO32- ions produce a white precipitate when the soluble carbonate salts is added to any metallic cation.

 

Procedure

Place about 2cm3 of Ca(NO3)2, Ba(NO3)2, MgCl2 ,Pb(NO3)2 andZnSO4 in separate boiling tubes.

 

Add six drops of Potassium /sodium carbonate(IV)/ ammonium carbonate (IV)solution.

 

Observation Inference Explanation
Green precipitate Cu2+ ,Fe2+,ions Copper(II)carbonate(IV) and Iron(II)
carbonate (IV) are precipitated as insoluble
CO3 2-(aq) green precipitates.
Cu2+(aq)+ CO3 2-(aq) -> CuCO3(s)
Fe2+(aq)+ CO32-(aq) -> FeCO3(s)
When sodium carbonate(IV)is added to
CuCO3(s) the CO3 2-(aq) ions are first
hydrolysed to produce CO2(g) and OH
(aq)ions.
CO3 2-(aq) + H2O (l) -> CO2 (g) + 2OH (aq)
The OH(aq) ions further react to form basic
copper(II) carbonate(IV). Basic copper(II)
carbonate(IV) is the only green salt of copper.
Cu2+(aq)+ CO3 2-(aq)+2OH (aq)

 

 

->CuCO3.Cu(OH)2 (s)
White precipitate CO3 2- White ppt of the carbonate(IV)salt is
precipitated
Ca2+(aq) + CO3 2- (aq)  ->  CaCO3(s)
Mg2+(aq) + CO3 2- (aq)  ->  MgCO3(s)
Pb2+(aq) + CO3 2- (aq)  ->  PbCO3(s)
Zn2+(aq) + CO3 2- (aq)   ->  ZnCO3(s)

Note

 

(i)Iron(III)carbonate(IV) does not exist. (ii)Copper(II)Carbonate(IV) exist only as the basic CuCO3.Cu(OH) 2 (iii)Both BaCO3 and BaSO3 are insoluble white precipitate. If hydrochloric

acid is added to the white precipitate;

 

  1. BaCO3 produces CO2 gas. When bubbled/directed into lime water solution,a white precipitate is formed.

 

  1. BaSO3 produces SO2 gas. When bubbled/directed into orange acidified potassium dichromate(VI) solution, it turns to green/decolorizes acidified potassium manganate(VII).

 

  • Reaction of cation with sulphide / S2- ions

All sulphides are insoluble black solids/precipitates except sodium sulphide/ Na2S/ potassium sulphide/K2S.When a few/3drops of the soluble sulphide is added to a metal cation/salt, a black precipitate is formed.

 

Procedure

Place about 2cm3 of Cu(NO3)2, FeSO4, MgCl2,Pb(NO3)2 and ZnSO4 in separate boiling tubes.

Add six drops of Potassium /sodium sulphide solution.

 

Observation Inference Explanation
Black ppt S2- ions CuS, FeS,MgS,PbS, ZnS are black insoluble
precipitates
Cu2+(aq) + S2-(aq)  -> CuS(s)
Pb2+(aq) + S2-(aq)  -> PbS(s)
Fe2+(aq) + S2-(aq)  ->  FeS(s)
Zn2+(aq)  +  S2-(aq)  ->  ZnS(s)
Sample qualitative analysis guide

 

 

You are provided with solid Y(aluminium (III)sulphate(VI)hexahydrate).Carry out the following tests and record your observations and inferences in the space provided.

 

1(a) Appearance
Observations inference (1mark)
White crystalline solid Coloured ions Cu2+ , Fe2+ ,Fe3+ absent

 

 

(b)Place about a half spatula full of the solid into a clean dry boiling tube. Heat gently then strongly.

 

 

Observations

 

inference

 

(1mark)

 

 

Colourless droplets formed on the cooler

part of the test tube

Solid remains a white residue

 

 

 

 

Hydrated compound/compound containing water of crystallization

 

 

 

(c)Place all the remaining portion of the solid in a test tube .Add about 10cm3 of distilled water. Shake thoroughly. Divide the mixture into five portions.

 

Observation Inference    (1mark)
Solid dissolves to form Polar soluble compound
a colourless solution Cu2+ , Fe2+ ,Fe3+ absent

 

 

(i)To the first portion, add three drops of sodium hydroxide then add excess of the alkali.

 

Observation Inference    (1mark)
White ppt, soluble in excess Zn2+ , Pb2+ , Al3+

 

(ii)To the second portion, add three drops of aqueous ammonia then add excess of the alkali.

 

Observation Inference    (1mark)
White ppt, insoluble in excess Pb2+ , Al3+

 

(iii)To the third portion, add three drops of sodium sulphate(VI)solution.

 

Observation Inference    (1mark)
No white ppt Al3+

 

 

(iv)I.To the fourth portion, add three drops of Lead(II)nitrate(IV)solution. Preserve

 

Observation Inference    (1mark)
White ppt CO3 2-, SO4 2-, SO3 2-, Cl,

 

II.To the portion in (iv) I above , add five drops of dilute hydrochloric acid.

 

 

Observation

 

 

Inference

 

 

(1mark)

 

 

White ppt persist/remains

 

 

SO4

 

2-,  Cl-,

 

 

 

III.To the portion in (iv) II above, heat to boil.

 

 

Observation

 

 

Inference

 

 

(1mark)

 

 

White ppt persist/remains

 

 

SO4

 

2-,

 

 

 

 

Note that:

 

(i)From test above, it can be deduced that solid Y is hydrated aluminium(III)sulphate(VI) solid

 

(ii)Any ion inferred from an observation below must be derived from previous correct observation and inferences above. e.g.

 

Al3+ in c(iii) must be correctly inferred in either/or in c(ii) or c(i)above SO42- in c(iv)III must be correctly inferred in either/or in c(iv)II or c(iv)I

above

 

(iii)Contradiction in observations and inferences should be avoided.e.g. White ppt soluble in excessto infer presence of Al3+ ,Ba2+ ,Pb3+

 

(iv)Symbols of elements/ions should be correctly capitalized. e.g.

SO4-2is wrong,sO42-is wrong,cu2+is wrong.

 

Sample solutions of salt were labeled as I,II, III and IV. The actual solutions, not in that order are lead nitrate, zinc sulphate potassium chloride and calcium chloride.

 

a)When aqueous sodium carbonate was added to each sample separately, a white precipitate was formed in I, III and IV only. Identify solution II.

 

 

b)When excess sodium hydroxide was added to each sample separately, a white precipitate was formed in solutions III and I only.

 

Identify solution I

 

 

 

17.When solids/salts /solutes are added to a solvent ,some dissolve to form a solution.

Solute               +            Solvent            ->          Solvent

 

If a solution has a lot of solute dissolved in a solvent ,it is said to be concentrated.

 

If a solution has little solute dissolved in a solvent ,it is said to be dilute.

 

There is a limit to how much solute can dissolve in a given /specified amount of solvent/water at a given /specified temperature.

 

The maximum mass of salt/solid/solute that dissolve in 100g of solvent/water at a specified temperature is called solubility of a salt.

 

When no more solute can dissolve in a given amount of solvent at a specified temperature, a saturated solution is formed.

 

For some salts, on heating, more of the salt/solid/solute dissolve in the saturated solution to form a super saturated solution.

 

The solubility of a salt is thus calculated from the formula

 

Solubility = Mass of solute/salt/solid      x 100

Mass/volume of water/solvent

 

Practice examples

 

(a)Calculate the solubility of potassium nitrate(V) if 5.0 g of the salt is dissolved in 50.0cm3 of water.

 

Solubility = Mass of solute/salt/solid x 100 =>( 5.0 x 100 )  = 10.0 g /100g H 2 O
Mass/volume of water/solvent 50.0

 

(b)Calculate the solubility of potassium chlorate(V) if 50.0 g of the salt is dissolved in 250.0cm3 of water.

 

Solubility = Mass of solute/salt/solid x 100   =>( 50.0 x 100 )    = 20.0 g /100g H2O

 

 

Mass/volume of water/solvent 250.0

 

(c)If the solubility of potassium chlorate(V) is 5g/100g H2O at 80oC,how much can dissolve in 5cm3 of water at 80oC .

 

Mass of solute/salt/solid = Solubility x Mass/volume of water/solvent 100

 

=>            5 x 5 = 0.25g of KClO3 dissolve 100

 

 

(d)If the solubility of potassium chlorate(V) is 72g/100g H2O at 20oC,how much can saturate 25g of water at 20oC .

 

Mass of solute/salt/solid = Solubility x Mass/volume of water/solvent 100

 

=>               72 x 25 = 18.0g of KClO3 dissolve/saturate 100

 

 

  • 22g of potassium nitrate(V) was dissolved in 40.0g of water at 10o Calculate the solubility of potassium nitrate(V) at 10oC.

 

Solubility = Mass of solute/salt/solid x 100 =>( 22 x 100 )  = 55.0 g /100g H 2 O
Mass/volume of water/solvent 40.0.

 

(f)What volume of water should be added to 22.0g of water at 10oC if the solubility of KNO3 at 10oC is 5.0g/100g H2O?

 

Solubility is mass/100g H2O => 22.0g + x = 100cm3/100g H2O

X= 100 22 = 78 cm3 of H2O

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

  1. A graph of solubility against temperature is called solubility curve. It shows the influence of temperature on solubility of different substances/solids/salts.

 

Some substances dissolve more with increase in temperature while for others dissolve less with increase in temperature

 

 

 

 

 

KNO3

 

KClO3

 

 

Saturated solution of KClO3

 

Solubility

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Note:

(i)solubility of KNO3 and KClO3 increase with increase in temperature.

 

(ii)solubility of KNO3 is always higher than that of KClO3 at any specified temperature.

 

(iii)solubility of NaCl decrease with increase in temperature.

(iv)NaCl has the highest solubility at low temperature while KClO3 has the lowest solubility at low temperature.

 

(v)At point A both NaCl and KNO3 are equally soluble.

 

(vi)At point B both NaCl and KClO3 are equally soluble.

 

  • An area above the solubility curve of the salt shows a saturated /supersaturated solution.

 

  • An area below the solubility curve of the salt shows an unsaturated

 

19.(a) For salts whose solubility increases with increase in temperature, crystals form when the salt solution at higher temperatures is cooled to a lower temperature.

 

 

  • For salts whose solubility decreases form when the salt solution at lower

 

 

with increase in temperature, crystals temperatures is heated to a higher

 

 

The examples below shows determination of the mass of crystals deposited with changes in temperature.

 

1.The solubility of KClO3 at 100oC is 60g/100g water .What mass of KClO3 will be deposited at:

 

(i)75 oC if the solubility is now 39g/100g water.

At 100oC =  60.0g
Less at 75oC = – 39.0g
Mass of crystallized out 21.0g

 

(i)35 oC if the solubility is now 28 g/100g water.

At 100oC =  60.0g
Less at 35oC = – 28.0.0g
Mass of crystallized out 32.0g

 

  1. KNO3 has a solubility of 42 g/100g water at 20oThe salt was heated and

added 38g more of the solute which dissolved at100oC. Calculate the solubility of KNO3 at 100oC.

 

Solubility of KNO3 at 100oC = solubility at 20oC + mass of KNO3 added => 42g + 38g = 80g KNO3 /100g H2O

 

  1. A salt solution has a mass of 65g containing 5g of solute. The solubility of this salt is 25g per 100g water at 20o 60g of the salt are added to the solution at 20oC.Calculate the mass of the solute that remain undissolved.

 

Mass of solvent at 20oC = mass of solution mass of solute

=> 65   – 5 = 60g
Solubility before adding salt = mass of solute x 100
Volume of solvent
=> 5 x 100 = 8.3333g/100g water
60
Mass of solute to equalize with solubility = 25 8.3333g  = 16.6667g
Mass of solute undissolved =  60.0 – 16.6667g  = 43.3333 g
4. Study the table below
Salt Solubility in gram at

 

 

50oC 20oC
KNO3 90 30
KClO3 20 6

 

(i)What happens when the two salts are dissolved in water then cooled from 50oC to 20oC.

 

(90 30) = 60.0 g of KNO3 crystals precipitate

(20 6) = 14.0 g of KClO3 crystals precipitate

 

(ii)State the assumption made in (i) above.

 

Solubility of one salt has no effect on the solubility of the other.

 

  1. 5.0 g of hydrated potassium carbonate (IV) K2CO3.xH2O on heating leave

7.93 of the hydrate.

 

(a)Calculate the mass of anhydrous salt obtained.

Hydrated on heating leave anhydrous   = 7.93 g

 

(b)Calculate the mass of water of crystallization in the hydrated salt

 

Mass of water of crystallization = hydrated anhydrous => 10.0 – 7.93 = 2.07 g

 

(c)How many moles of anhydrous salt are there in 10of hydrate? (K= 39.0,C=12.0.O= 16.0)

 

Molar mass K2CO3= 138
Moles K2CO3 = mass of K 2 CO 3=> 7.93 = 0.0515 moles
Molar mass K2CO3 138

 

(d)How many moles of water are present in the hydrate for every one mole of K2CO3 ? (H=1.0.O= 16.0)

 

Molar mass H2O = 18
Moles H2O  = mass of H2O => 2.07 = 0.115 moles
Molar mass H2O 18
Mole ratio H2O : K2CO3= 0.115 moles 2 = 2
0.0515 moles 1

 

(e)What is the formula of the hydrated salt?

 

K2CO3 .2 H2O

 

  1. The table below shows the solubility of Potassium nitrate(V) at different temperatures.

 

Temperature(oC) 5.0 10.0 15.0 30.0 40.0 50.0 60.0
mass KNO3/ 15.0 20.0 25.0 50.0 65.0 90.0 120.0
100g water

 

(a)Plot a graph of mass of  in 100g water(y-axis) against temperature in oC

 

 

 

 

 

 

(b)From the graph show and determine (i)the mass of KNO3 dissolved at:

  1. 20oC

From a correctly plotted graph = 32g

  1. 35oC

From a correctly plotted graph = 57g

III. 55oC

From a correctly plotted graph = 104g

(ii)the temperature at which the following mass of KNO3 dissolved:

  1. 22g

From a correctly plotted graph =13.0oC

  1. 30g

From a correctly plotted graph =17.5oC

III.100g

From a correctly plotted graph =54.5oC

 

(c)Explain the shape of your graph.

Solubility of KNO3 increase with increase in temperature/More KNO3 dissolve as temperature rises.

 

(d)Show on the graph the supersaturated and unsaturated solutions. Above the solubility curve write; supersaturated

 

Below the solubility curve write; unsaturated

 

(e)From your graph, calculate the amount of crystals obtained when a saturated solution of KNO3 containing 180g of the salt is cooled from 80oC to:

 

I. 20oC
Solubility before heating = 180 g
Less  Solubility after heating(from the graph) = 32 g
Mass of KNO3crystals = 148 g
II. 35oC
Solubility before heating = 180 g
Less  Solubility after heating(from the graph) = 58 g
Mass of KNO3crystals = 122 g
III. 55oC

 

 

Solubility before heating = 180 g
Less  Solubility after heating(from the graph) = 102 g
Mass of KNO3crystals = 78 g

 

  1. The table below shows the solubility of salts A and B at various temperatures.

 

Temperature(oC) 0.0 10.0 20.0 30.0 40.0 50.0 60.0 70.0 80.0
Solubility of A 28.0 31.0 34.0 37.0 40.0 43.0 45.0 48.0 51.0
Solubility of B 13.0 21.0 32.0 46.0 64.0 85.0 110.0 138.0 169.0

 

(a)On the same axis plot a graph of solubility (y-axis) against temperature for each salt.

(b)At what temperature are the two salts equally soluble.

The point of intersection of the two curves = 24oC

 

(c)What happens when a mixture of 100g of salt B with 100g if water is heated to 80oC

From the graph, the solubility of B at 80oC is 169g /100g water. All the 100g crystals of B dissolve.

 

(d)What happens when the mixture in (c) above is then cooled from 50oC to

20oC.
Method I.
Total mass before cooling at 50oC = 100.0 g
(From graph) Solubility/mass after cooling at 20oC = 32.0 g
Mass of crystals deposited 68.0 g
Method II.
Mass of soluble salt crystals at 50oC added = 100 g
(From graph)Solubility/mass before cooling at 50oC = 85.0 g
Mass of crystals that cannot dissolve at 50oC 15.0  g
(From graph) Solubility/mass before cooling at 50oC = 85.0 g
(From graph) Solubility/mass after cooling at 20oC = 32.0 g
Mass of crystals deposited after cooling 53.0 g
Total mass of crystals deposited = 15.0 + 53.0 = 68.0 g

 

 

(e)A mixture of 40g of A and 60g of B is added to 10g of water and heated to 70oC.The solution is then allowed to cool to 10oC.Describe clearly what happens.

 

I.For salt A
Solubility of A before heating = mass of A  x 100
Volume of water added
=> 40   x  100= 400g/100g Water
10
(Theoretical)Solubility of A before heating = 400 g
Less (From graph ) Solubility of A after heating at 70oC = 48g
Mass of crystals that can not dissolve at70oC = 352 g
(From graph ) Solubility of A after heating at 70oC = 48g
Less (From graph ) Solubility of A after cooling to 10oC = 31g

 

Mass of crystals that crystallize out on cooling to10oC = 17 g
Mass of crystals that can not dissolve at70oC = 352 g
Add Mass of crystals that crystallize out on cooling to10oC = 17 g
Total mass of A that does not dissolve/crystallize/precipitate = 369 g
I.For salt B
Solubility of B before heating = mass of B  x 100
Volume of water added
=> 60   x  100=   600g/100g Water
10
(Theoretical)Solubility of B before heating = 600 g
Less (From graph ) Solubility of B after heating at 70oC = 138g
Mass of crystals that cannot dissolve at70oC = 462 g
(From graph ) Solubility of B after heating at 70oC = 138g
Less (From graph ) Solubility of B after cooling to 10oC = 21g
Mass of crystals that crystallize out on cooling to10oC = 117 g
Mass of crystals that cannot dissolve at70oC = 462 g
Add Mass of crystals that crystallize out on cooling to10oC = 117 g
Total mass of A that does not dissolve/crystallize/precipitate = 579 g

 

(f)State the assumption made in (e)above

Solubility of one salt has no effect on the solubility of the other

 

  1. When 5.0 g of potassium chlorate (V) was put in 10cm3 of water and heated, the solid dissolves. When the solution was cooled , the temperature at which crystals reappear was noted. Another 10cm3 of water was added and the mixture heated to dissolve then cooled for the crystals to reappear .The table below shows the the results obtained

 

Total volume of water added(cm3) 10.0 20.0 30.0 40.0 50.0
Mass of KClO3 5.0 5.0 5.0 5.0 5.0
Temperature at which crystals appear 80.0 65.0 55.0 45.0 30.0
Solubility of KclO3 50.0 25.0 16.6667 12.5 10.0

 

(a)Complete the table to show the solubility of KclO3 at different temperatures.

 

(b)Plot a graph of mass of KClO3 per 100g water against temperature at which crystals form.

 

(c)From the graph, show and determine ;

(i)the solubility of KClO3 at

  1. 50oC

From a well plotted graph = 14.5 g KClO3/100g water

  1. 35oC

 

From a well plotted graph = 9.0 g KclO3/100g water

(ii)the temperature at which the solubility is:

 

I.10g/100g water

From a well plotted graph = 38.0 oC

II.45g/100g water

From a well plotted graph = 77.5 oC

 

(d)Explain the shape of the graph.

Solubility of KClO3 increase with increase in temperature/more KclO3dissolve as temperature rises.

 

(e)What happens when 100g per 100g water is cooled to 35.0 oC

Solubility before heating = 100.0

 

(From the graph)  Solubility after cooling   = 9.0
Mass of salt precipitated/crystallization = 91.0 g

 

  1. 0cm3 of water dissolved various masses of ammonium chloride crystals at different temperatures as shown in the table below.

 

Mass of ammonium chloride(grams) 4.0 4.5 5.5 6.5 9.0
Temperature at which solid dissolved(oC) 30.0 50.0 70.0 90.0 120.0
Solubility of NH4Cl 16.0 18.0 22.0 26.0 36.0
(a)Complete the table
(b)Plot a solubility curve

 

(c)What happens when a saturated solution of ammonium chloride is cooled from 80oC to 40oC.

 

(From the graph )Solubility at 80oC = 24.0 g
Less  (From the graph )Solubility at 40oC = 16.8 g
Mass of crystallized/precipitated =  7.2 g

 

  1. Solubility and solubility curves are therefore used
    • to know the effect of temperature on the solubility of a salt

 

(ii)to fractional crystallize two soluble salts by applying their differences in solubility at different temperatures.

(iii)determine the mass of crystal that is obtained from crystallization.

 

21.Natural fractional crystallization takes place in Kenya/East Africa at:

 

  • Lake Magadi during extraction of soda ash(Sodium carbonate) from Trona(sodium sesquicarbonate)

 

  • Ngomeni near Malindi at the Indian Ocean Coastline during the extraction of common salt(sodium chloride).

 

22.Extraction of soda ash from Lake Magadi in Kenya

 

Rain water drains underground in the great rift valley and percolate underground where it is heated geothermically.

 

 

The hot water dissolves underground soluble sodium compounds and comes out on the surface as alkaline springs which are found around the edges of Lake Magadi in Kenya.

Temperatures around the lake are very high (30-40oC) during the day.

 

The solubility of trona decrease with increase in temperature therefore solid crystals of trona grows on top of the lake (upto or more than 30metres thick)

 

A bucket dredger mines the trona which is then crushed ,mixed with lake liquor and pumped to washery plant where it is further refined to a green granular product called CRS.

 

The CRS is then heated to chemically decompose trona to soda ash(Sodium carbonate)

 

Chemical equation

2Na2CO3.NaHCO3.2H2O(s) -> 3Na2CO3 (s) + CO2(g) + 5H2O(l)

 

Soda ash(Sodium carbonate) is then stored .It is called Magadi Soda. Magadi Soda is used :

  • make glass
  • for making soapless detergents
  • softening hard water.

(iv)

 

Common salt is colledcted at night because its solubility decreases with decrease in temperature. It is used as salt lick/feed for animals.

 

Summary flow diagram showing the extraction of Soda ash  from Trona

 

 

Sodium chloride and Trona

 

dissolved in the sea

 

Carbon(IV) oxide

 

 

Natural fractional crystallization

 

 

 

Crystals of Trona Dredging Crushing Furnace
(Day time) /scooping/ (Heating)
digging

 

Crystals of sodium

Soda ash

chloride(At night)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

23.Extraction of common salt from Indian Ocean at Ngomeni in Kenya

 

Oceans are salty.They contain a variety of dissolved salts (about 77% being sodium chloride).

 

During high tide ,water is collected into shallow pods and allowed to crystallize as evaporation takes place.The pods are constructed in series to increase the rate of evaporation.

 

At the final pod ,the crystals are scapped together,piled in a heap and washed with brine (concentrated sodium chloride).

 

It contains MgCl2 and CaCl2 . MgCl2 and CaCl2are hygroscopic. They absorb water from the atmosphere and form a solution.

 

This makes table salt damp/wet on exposure to the atmosphere.

 

24.Some water form lather easily with soap while others do not.

 

Water which form lather easily with soap is said to be “soft”

Water which do not form lather easily with soap is said to be “hard”

 

Hardness of water is caused by the presence of Ca2+ and Mg2+ ions. Ca2+ and Mg2+ ions react with soap to form an insoluble grey /white suspension/precipitate called Scum/ curd. Ca2+ and Mg2+ ions in water come from the water sources passing through rocks containing soluble salts of Ca2+ and Mg2+ e.g. Limestone or gypsum

 

There are two types of water hardness:

(a)temporary hardness of water

(b)permanent hardness of water

 

(a)temporary hardness of water

 

 

Temporary hardness of water is caused by the presence of dissolved calcium hydrogen carbonate/Ca(HCO3)2 and magnesium hydrogen carbonate/Mg(HCO3)2 When rain water dissolve carbon(IV) oxide from the air it forms waek carbonic(IV) acid i.e.

CO2(g)      +         H2O(l)    ->         H2CO3(aq)

 

When carbonic(IV) acid passes through limestone/dolomite rocks it reacts to form soluble salts i.e.

 

 

In limestone areas; In dolomite areas;

 

H2CO3(aq)

 

H2CO3(aq)

 

+   CaCO3(s)

+    MgCO3(s)

 

->

->

 

Ca(HCO3)2 (aq)

Mg(HCO3)2 (aq)

 

 

 

(b)permanent hardness of water

 

Permanent hardness of water is caused by the presence of dissolved calcium sulphate(VI)/CaSO4 and magnesium sulphate(VI)/Mg SO4 Permanent hardness of water is caused by water dissolving CaSO4 and MgSO4 from ground rocks.

 

Hardness of water can be removed by the following methods:

 

(a)Removing temporary hardness of water

 

(i)Boiling/heating.

 

Boiling decomposes insoluble calcium hydrogen carbonate/Ca(HCO3)2 and magnesium hydrogen carbonate/Mg(HCO3)2 to insoluble CaCO3 and MgCO3 that precipitate away. i.e

 

Chemical equation

Ca(HCO3)2(aq)    ->   CaCO3 (s)   +   CO2(g)   +     H2O(l)

Mg(HCO3)2(aq)    ->   MgCO3 (s)   +   CO2(g)   +    H2O(l)

 

(ii)Adding sodium carbonate (IV) /Washing soda.

 

Since boiling is expensive on a large scale ,a calculated amount of sodium

carbonate decahydrate /Na2CO3.10H2O  precipitates insoluble Ca2+(aq) and

Mg2+(aq) ions as carbonates to remove both temporary and permanent hardness of

 

water .This a double decomposition reaction where two soluble salts form an

insoluble and soluble salt. i.e.

 

(i)with temporary hard water

Chemical equation

Na2CO3  (aq)  +  Ca(HCO3) 2 (aq)  ->    NaHCO3(aq)  + CaCO3 (s)

Na2CO3  (aq)  +  Mg(HCO3) 2 (aq)  ->     NaHCO3(aq) + MgCO3 (s)

Ionic equation

CO32-  (aq)  +          Ca2+  (aq)      ->       CaCO3 (s)

 

CO32-  (aq)   +          Mg2+  (aq)     ->       MgCO3 (s)

 

(ii)with permanent hard water

Chemical equation

 

 

Na2CO3

 

Na2CO3

 

(aq)

 

(aq)

 

 

+ MgSO4 + CaSO4

 

(aq)

(aq)

 

->

->

 

 

Na2SO4 Na2SO4

 

 

(aq) + MgCO3

 

(aq) + MgCO3

 

(s)

(s)

 

 

Ionic equation

CO32-  (aq)  +          Ca2+  (aq)      ->       CaCO3 (s)

CO32-  (aq)  +           Mg2+  (aq)     ->       MgCO3 (s)

 

(iii)Adding calcium (II)hydroxide/Lime water

 

Lime water/calcium hydroxide removes only temporary hardness of water from by precipitating insoluble calcium carbonate(IV).

Chemical equation

Ca(OH)2  (aq)  +  Ca(HCO3) 2 (aq)  ->    2H2O(l)  + 2CaCO3 (s)

 

Excess of Lime water/calcium hydroxide should not be used because it dissolves again to form soluble calcium hydrogen carbonate(IV) causing the hardness again.

 

(iv)Adding aqueous ammonia

 

Aqueous ammonia removes temporary hardness of water by precipitating insoluble calcium carbonate(IV) and magnesium carbonate(IV)

 

Chemical equation

2NH3  (aq)  +  Ca(HCO3) 2 (aq)  ->     (NH4) 2CO3(aq)  + CaCO3 (s)

2NH3  (aq)  +  Mg(HCO3) 2 (aq)  ->     (NH4) 2CO3(aq)  + MgCO3 (s)

(v)Use of ion-exchange permutit

 

This method involves packing a chamber with a resin made of insoluble complex of sodium salt called sodium permutit.

 

The sodium permutit releases sodium ions that are exchanged with Mg2+ and Ca2+ ions in hard water making the water to be soft. i.e.

 

Na2X(aq) + Ca2+ (aq)    ->   Na+ (aq)  +   CaX(s)

 

 

Hard water containing Mg2+ and Ca2+

 

 

 

 

 

 

 

Ion exchange resin as

Sodium permutit

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Na+ ions replace Mg2+
——-
and Ca2+ to make the water soft.

 

 

 

 

 

 

 

 

When all the Na+ ions in the resin is fully exchanged with Ca2+ and Ng2+ ions in the permutit column ,it is said to be exhausted.

 

Brine /concentrated sodium chloride solution is passed through the permutit column to regenerated /recharge the column again.

 

Hard water containing Mg2+ and Ca2+

 

 

 

 

 

 

 

 

 

Ion exchange resin as

 

Sodium permutit

 

 

 

 

 

 

 

 

 

 

 

Na+ ions replace Mg2+
——-
and Ca2+ to make the water soft.

 

(vi)Deionization /demineralization

 

This is an advanced ion exchange method of producing deionized water .Deionized water is extremely pure water made only of hydrogen and oxygen only without any dissolved substances.

 

Deionization involve using the resins that remove all the cations by using:

 

(i)A cation exchanger which remove /absorb all the cations present in water and leave only H+ ions.

 

(ii)An anion exchanger which remove /absorb all the anions present in water and leave only OH ions.

The H+(aq) and OH (aq) neutralize each other to form pure water.

 

Chemical equation

H+(aq)  +   OH (aq)  ->       H2O(l)

 

When exhausted the cation exchanger is regenerated by adding H+(aq) from sulphuric(VI)acid/hydrochloric acid.

 

When exhausted the anion exchanger is regenerated by adding OH(aq) from sodium hydroxide.

 

Advantages of hard water

Hard water has the following advantages:

 

(i)Ca2+(aq) in hard water are useful in bone and teeth formation (ii) is good for brewing beer

 

(iii)contains minerals that cause it to have better /sweet taste (iv)animals like snails and coral polyps use calcium to make their shells

and coral reefs respectively.

 

(v)processing mineral water

 

Disadvantages of hard water

 

Hardness of water:

(i)waste a lot of soap during washing before lather is formed.

(ii)causes stains/blemishes/marks on clothes/garments

 

(iii)causes fur on electric appliances like kettle ,boilers and pipes form decomposition of carbonates on heating .This reduces their efficiency hence more/higher cost of power/electricity.

 

 

 

 

 

 

Sample revision questions

 

In an experiment, soap solution was added to three separate samples of water. The table below shows the volumes of soap solution required to form lather with 1000cm3 of each sample of water before and after boiling.

 

Sample Sample Sample
I II III
Volume of soap before water is boiled 27.0 3.0 10.0
(cm3)
Volume of soap after water is boiled(cm3) 27.0 3.0 3.0

 

 

  1. Which water sample is likely to be soft? Explain. (2mks)

 

  1. Name the change in the volume of soap solution used in sample III

 

(1mk)

On heating the sample water become soft bcause it is temporary hard.

 

2.Study the scheme below and use it to aanswer the questions that follow:

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(a)Write the formula of:

(i)Cation in solution K

Al3+

(ii)white ppt  L

Al(OH)3

 

  • colourless solution M [Al(OH)4]

 

  • colourless solution N AlCl3

 

(v)white ppt  P

Al(OH)3

 

(b)Write the ionic equation for the reaction for the formation of:

(i)white ppt  L

Al3+(aq)       + 3OH (aq)     ->     Al(OH)3(s)

 

(v)white ppt  P

Al3+(aq)       + 3OH (aq)     ->     Al(OH)3(s)

 

(c)What property is illustrated in the formation of colourless solution M and N.

 

Amphotellic

 

 

 

19.0.0 ENERGY CHANGES IN CHEMICAL AND PHYSICAL PROCESSES

 

(25 LESSONS)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

1.Introduction to Energy changes

 

Energy is the capacity to do work. There are many/various forms of energy like heat, electric, mechanical, and/ or chemical energy.There are two types of energy:

 

(i)Kinetic Energy(KE) ;the energy in motion.

(ii)Potential Energy(PE); the stored/internal energy.

 

Energy like matter , is neither created nor destroyed but can be transformed /changed from one form to the other/ is interconvertible. This is the principle of conservation of energy. e.g. Electrical energy into heat through a filament in bulb. Chemical and physical processes take place with absorption or evolution/production of energy mainly in form of heat

 

The study of energy changes that accompany physical/chemical reaction/changes is called Thermochemistry. Physical/chemical reaction/changes that involve energy changes are called thermochemical reactions. The SI unit of energy is the Joule(J).Kilo Joules(kJ)and megaJoules(MJ) are also used. The Joule(J) is defined as the:

 

  • quantity of energy transferred when a force of one newton acts through a distance of one metre.

 

  • quantity of energy transferred when one coulomb of electric charge is passed through a potential difference of one volt.

 

All thermochemical reactions should be carried out at standard conditions of:

  • 298K /25oC temperature

(ii)101300Pa/101300N/m2 /760mmHg/1 atmosphere pressure.

 

2.Exothermic and endothermic processes/reactions

 

Some reactions / processes take place with evolution/production of energy. They are said to be exothermic while others take place with absorption of energy. They are said to be endothermic.

 

Practically exothermic reactions / processes cause a rise in temperature (by a rise in thermometer reading/mercury or alcohol level rise)

 

Practically endothermic reactions / processes cause a fall in temperature (by a fall in thermometer reading/mercury or alcohol level decrease)

 

To demonstrate/illustrate exothermic and endothermic processes/reactions

 

  1. Dissolving Potassium nitrate(V)/ammonium chloride crystals

 

Procedure:

 

Measure 20cm3 of water in a beaker. Determine and record its temperature T1.Put about 1.0g of Potassium nitrate(V) crystals into the beaker. Stir the mixture carefully and note the highest temperature rise /fall T2.Repeat the

whole procedure by using ammonium chloride in place of Potassium nitrate

 

  • Sample results

 

Temperture (oC) Using Potassium Using Ammonium
nitrate(V) crystals chloride crystals
T2(Final temperature) 21.0 23.0
T1 (Initial temperature) 25.0 26.0
Change in temperature(T2 T1) 4.0 3.0

Note:

(i)Initial(T1) temperature of dissolution of both potassium nitrate(V) crystals and ammonium chloride crystals is higher than the final temperature(T2)

 

(ii) Change in temperature(T2 T1) is not a mathematical 4.0 or 3.0. (iii)Dissolution of both potassium nitrate(V) and ammonium chloride crystals is an endothermic process because initial(T1) temperature is higher than the final temperature(T2) thus causes a fall/drop in temperature.

 

  1. Dissolving concentrated sulphuric(VI) acid/sodium hydroxide crystals

 

2

 

Procedure:

 

Measure 20cm3 of water in a beaker. Determine and record its temperature T1.Carefully put about 1.0g/four pellets of sodium hydroxide crystals into the beaker. Stir the mixture carefully and note the highest temperature rise /fall T2.Repeat the whole procedure by using 2cm3 of concentrated sulphuric(VI) acid in place of sodium hydroxide crystals.

 

CAUTION:

 

(i)Sodium hydroxide crystals are caustic and cause painful blisters on contact with skin.

 

  • Concentrated sulphuric (VI) acid is corrosive and cause painful wounds on contact with skin.

 

Sample results

Temperture (oC) Using Sodium Using Concentrated
hydroxide pellets sulphuric(VI) acid
T2(Final temperature) 30.0 32.0
T1 (Initial temperature) 24.0 25.0
Change in temperature(T2 T1) 6.0 7.0

 

Note:

(i)Initial (T1) temperature of dissolution of both concentrated sulphuric (VI) acid and sodium hydroxide pellets is lower than the final temperature (T2). (ii)Dissolution of both Sodium hydroxide pellets and concentrated sulphuric (VI) acid is an exothermic process because final (T2) temperature is higher than the initial temperature (T1) thus causes a rise in temperature.

 

The above reactions show heat loss to and heat gain from the surrounding as illustrated by a rise and fall in temperature/thermometer readings.

 

Dissolving both potassium nitrate(V) and ammonium chloride crystals causes heat gain from the surrounding that causes fall in thermometer reading.

 

Dissolving both Sodium hydroxide pellets and concentrated sulphuric (VI) acid causes heat loss to the surrounding that causes rise in thermometer reading.

 

At the same temperature and pressure ,heat absorbed and released is called enthalpy/ heat content denoted H.

 

Energy change is measured from the heat content/enthalpy of the final and initial products. It is denoted H(delta H).i.e.

Enthalpy/energy/ change in heat content ∆H = HfinalHinitial

 

For chemical reactions:

H = HproductsHreactants

 

For exothermic reactions, the heat contents of the reactants is more than/higher than the heat contents of products, therefore the H is negative (∆H)

 

3

 

For endothermic reactions, the heat contents of the reactants is less than/lower than the heat contents of products, therefore the H is negative (+∆H)

 

Graphically, in a sketch energy level diagram:

 

(i)For endothermic reactions the heat content of the reactants should be relatively/slightly lower than the heat content of the products

 

(ii)For exothermic reactions the heat content of the reactants should be relatively/slightly higher than the heat content of the products

 

Sketch energy level diagrams for endothermic dissolution

 

 

 

Energy

 

(kJ)

 

H2

 

KNO3(aq)

 

 

 

+H=H2H1

 

H1    KNO3(s)

 

 

 

Reaction path/coordinate

 

 

 

 

Energy
(kJ)H2 NH4Cl (aq)
+H=H2H1
H1 NH4Cl (s)

 

 

 

Reaction path/coordinate/progress Sketch energy level diagrams for exothermic dissolution

 

4

 

 

 

H2        NaOH (s)

 

Energy(kJ)

H=H2H1

 

H1                                                 NaOH (aq)

 

Reaction path/coordinate/progress

 

 

 

 

H2       H2SO4 (l)

Energy

(kJ)

H=H2H1

 

H1                                                 H2SO4 (aq)

 

Reaction path/coordinate/progress

 

 

 

 

 

 

 

 

 

 

 

 

 

 

3.Energy changes in physical processes

 

Melting/freezing/fusion/solidification and boiling/vaporization/evaporation are the two physical processes. Melting /freezing point of pure substances is fixed /constant. The boiling point of pure substance depend on external atmospheric pressure.

 

Melting/fusion is the physical change of a solid to liquid. Freezing is the physical change of a liquid to solid.

 

Melting/freezing/fusion/solidification are therefore two opposite but same reversible physical processes. i.e

 

A (s)  ========A(l)

 

5

 

Boiling/vaporization/evaporation is the physical change of a liquid to gas/vapour. Condensation/liquidification is the physical change of gas/vapour to liquid. Boiling/vaporization/evaporation and condensation/liquidification are therefore two opposite but same reversible physical processes. i.e

 

B (l)   ========B(g)

Practically

 

  • Melting/liquidification/fusion involves heating a solid to weaken the strong bonds holding the solid particles together. Solids are made up of very strong bonds holding the particles very close to each other (Kinetic Theory of matter).On heating these particles gain energy/heat from the surrounding heat source to form a

 

liquid with weaker bonds holding the particles close together but with some degree of freedom. Melting/freezing/fusion is an endothermic (+H)process that require/absorb energy from the surrounding.

 

(ii)Freezing/fusion/solidification involves cooling a a liquid to reform /rejoin the very strong bonds to hold the particles very close to each other as solid and thus lose their degree of freedom (Kinetic Theory of matter). Freezing /fusion / solidification is an exothermic (H)process that require particles holding the liquid together to lose energy to the surrounding.

 

(iii)Boiling/vaporization/evaporation involves heating a liquid to completely break/free the bonds holding the liquid particles together. Gaseous particles have high degree of freedom (Kinetic Theory of matter). Boiling /vaporization / evaporation is an endothermic (+H) process that require/absorb energy from the surrounding.

 

(iv)Condensation/liquidification is reverse process of boiling /vaporization / evaporation.It involves gaseous particles losing energy to the surrounding to form a liquid.It is an exothermic(+H) process.

 

The quantity of energy required to change one mole of a solid to liquid or to form one mole of a solid from liquid at constant temperature is called molar enthalpy/latent heat of fusion. e.g.

 

H2O(s)  -> H2O(l) H = +6.0kJ mole-1  (endothermic process)

H2O(l)  -> H2O(s) H = -6.0kJ mole-1 (exothermic process)

 

The quantity of energy required to change one mole of a liquid to gas/vapour or to form one mole of a liquid from gas/vapour at constant temperature is called molar enthalpy/latent heat of vapourization. e.g.

 

H2O(l)   -> H2O(g) H = +44.0kJ mole-1  (endothermic process)

H2O(g)  -> H2O(l) H = -44.0kJ mole-1 (exothermic process)

 

6

 

The following experiments illustrate/demonstrate practical determination of melting and boiling

 

  1. To determine the boiling point of water Procedure:

 

Measure 20cm3 of tap water into a 50cm3 glass beaker. Determine and record its temperature.Heat the water on a strong Bunsen burner flame and record its temperature after every thirty seconds for four minutes.

 

Sample results

Time(seconds) 0 30 60 90 120 150 180 210 240
Temperature(oC) 25.0 45.0 85.0 95.0 96.0 96.0 96.0 97.0 98.0

 

Questions

1.Plot a graph of temperature against time(y-axis)

 

Sketch graph of temperature against time

 

 

 

 

 

 

 

 

 

 

boiling point

 

96 oC

 

 

 

Temperature(0C)

 

 

 

25oC

 

time(seconds)

 

2.From the graph show and determine the boiling point of water

Note:

 

Water boils at 100oC at sea level/one atmosphere pressure/101300Pa but boils at below 100oC at higher altitudes. The sample results above are from Kiriari Girls High School-Embu County on the slopes of Mt Kenya in Kenya. Water here boils at 96oC.

 

7

 

3.Calculate the molar heat of vaporization of water.(H= 1.0,O= 16.O)

 

Working:

 

Mass of water = density x volume => (20 x 1) /1000 = 0.02kg Quantity of heat produced

 

  • mass of water x specific heat capacity of water x temperature change =>0.02kg x 4.2 x ( 96 25 ) = 964kJ

Heat of vaporization of one mole H2O =      Quantity of heat

Molar mass of H2O

=>5.964kJ  =         0.3313 kJ mole -1

18

 

To determine the melting point of candle wax Procedure

 

Weigh exactly 5.0 g of candle wax into a boiling tube. Heat it on a strongly Bunsen burner flame until it completely melts. Insert a thermometer and remove the boiling tube from the flame. Stir continuously. Determine and record the temperature after every 30seconds for four minutes.

Sample results

 

Time(seconds) 0 30 60 90 120 150 180 210 240
Temperature(oC) 93.0 85.0 78.0 70.0 69.0 69.0 69.0 67.0 65.0

 

Questions

 

1.Plot a graph of temperature against time(y-axis) Sketch graph of temperature against time

 

 

 

 

 

93 oC

 

 

 

 

Temperature(0C)

 

69oC

 

 

 

 

melting point

 

 

 

 

time(seconds)

 

2.From the graph show and determine the melting point of the candle wax

 

 

 

 

 

 

 

8

 

4.Energy changes in chemical processes

 

Thermochemical reactions measured at standard conditions of 298K(25oC) and 101300Pa/101300Nm2/ 1 atmospheres/760mmHg/76cmHg produce standard enthalpies denoted Hᶿ.

 

Thermochemical reactions are named from the type of reaction producing the energy change. Below are some thermochemical reactions:

  • Standard enthalpy/heat of reaction Hᶿr
  • Standard enthalpy/heat of combustion Hᶿc

 

  • Standard enthalpy/heat of displacement Hᶿd
  • Standard enthalpy/heat of neutralization Hᶿn

 

  • Standard enthalpy/heat of solution/dissolution Hᶿs

 

  • Standard enthalpy/heat of formation Hᶿf

 

(a)Standard enthalpy/heat of reaction ∆Hᶿr

 

The molar standard enthalpy/heat of reaction may be defined as the energy/heat change when one mole of products is formed at standard conditions

 

A chemical reaction involves the reactants forming products. For the reaction to take place the bonds holding the reactants must be broken so that new bonds of the products are formed. i.e.

 

AB + C-D -> A-C + B-D

 

Old Bonds broken A-B and C-D on reactants New Bonds formed A-C and B-D on products

 

The energy required to break one mole of a (covalent) bond is called bond dissociation energy. The SI unit of bond dissociation energy is kJmole-1 The higher the bond dissociation energy the stronger the (covalent)bond

 

Bond dissociation energies of some (covalent)bonds

Bond Bond dissociation energy Bond dissociation energy
(kJmole-1) (kJmole-1)
H-H 431 I-I 151
C-C 436 C-H 413
C=C 612 O-H 463
C = C 836 C-O 358
N = N 945 H-Cl 428
N-H 391 H-Br 366
F-F 158 C-Cl 346
Cl-Cl 239 C-Br 276
Br-Br 193 C-I 338
H-I 299 O=O 497
Si-Si 226 C-F 494
9

 

The molar enthalpy of reaction can be calculated from the bond dissociation energy by:

(i)adding the total bond dissociation energy of the reactants(endothermic

 

process/+∆H) and total bond dissociation energy of the products(exothermic process/∆H).

 

(ii)subtracting total bond dissociation energy of the reactants from the total bond dissociation energy of the products(exothermic process/∆H less/minus endothermic process/+∆H).

 

Practice examples/Calculating Hr 1.Calculate ∆Hr from the following reaction:

  1. H2(g) +    Cl2(g)   ->  2HCl(g)

 

Working

Old bonds broken (endothermic process/+∆H )

 

  • (H-H + Cl-Cl) => (+431 + (+ 239)) = + 670kJ New bonds broken (exothermic process/H )

 

  • (2(H-Cl ) => (- 428 x 2)) = -856kJ

 

Hr =( + 670kJ + -856kJ) =          186 kJ = -93kJ mole-1 2

 

The above reaction has negative H enthalpy change and is therefore practically exothermic.

The thermochemical reaction is thus: HCl(g) ∆Hr =  -93kJ
½ H2(g) +   ½ Cl2(g) ->
b) CH4(g) +    Cl2(g) -> CH3Cl + HCl(g)

Working

 

Old bonds broken (endothermic process/+∆H )

 

= (4(C-H) + Cl-Cl)

=> ((4 x +413) + (+ 239)) = + 1891kJ

New bonds broken (exothermic process/H )

 

= (3(C-H + H-Cl + C-Cl)

 

=> (( 3 x – 413) + 428 + 346) = –2013 kJ Hr =( + 1891kJ + -2013 kJ) = -122 kJ mole-1

 

The above reaction has negative H enthalpy change and is therefore practically exothermic.

The thermochemical reaction is thus: CH3Cl(g)  +  HCl(g)   ∆H = -122 kJ
CH4(g) + Cl2(g)   ->
c)  CH2CH2(g) + Cl2(g)   -> CH3Cl CH3Cl (g)
10

 

Working

Old bonds broken (endothermic process/+∆H )

 

= (4(C-H) + Cl-Cl + C=C)

 

=> ((4 x +413) + (+ 239) +(612)) = + 2503kJ New bonds broken (exothermic process/H )

= (4(C-H + C-C + 2(C-Cl) )

 

=> (( 3 x – 413) + -436 +2 x 346 = –2367 kJ Hr =( + 2503kJ + -2367 kJ) = +136 kJ mole-1

 

The above reaction has negative +H enthalpy change and is therefore practically endothermic.

The thermochemical reaction is thus:

CH2CH2(g)    +     Cl2(g)    ->   CH3Cl CH3Cl (g)       ∆H      = +136 kJ

Note that:

 

(i)a reaction is exothermic if the bond dissociation energy of reactants is more than bond dissociation energy of products.

 

(ii)a reaction is endothermic if the bond dissociation energy of reactants is less than bond dissociation energy of products.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

11

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(b)Standard enthalpy/heat of combustion ∆Hᶿc

 

The molar standard enthalpy/heat of combustion(Hᶿc) is defined as the energy/heat change when one mole of a substance is burnt in oxygen/excess air at standard conditions.

 

Burning is the reaction of a substance with oxygen/air. It is an exothermic process producing a lot of energy in form of heat.

 

A substance that undergoes burning is called a fuel. A fuel is defined as the combustible substance which burns in air to give heat energy for domestic or industrial use. A fuel may be solid (e.g coal, wood, charcoal) liquid (e.g petrol, paraffin, ethanol, kerosene) or gas (e.g liquefied petroleum gas/LPG, Water gas-CO2/H2, biogas-methane, Natural gas-mixture of hydrocarbons)

 

To determine the molar standard enthalpy/heat of combustion(∆Hᶿc) of ethanol

 

Procedure

 

Put 20cm3 of distilled water into a 50cm3 beaker. Clamp the beaker. Determine the temperature of the water T1.Weigh an empty burner(empty tin with wick).

 

12

 

Record its mass M1.Put some ethanol into the burner. Weigh again the burner with the ethanol and record its mass M2. Ignite the burner and place it below the clamped 50cm3 beaker. Heat the water in the beaker for about one minute. Put off the burner. Record the highest temperature rise of the water, T2. Weigh the burner again and record its mass M3

 

Sample results:

 

Volume of water used 20cm3
Temperature of the water before heating T1 25.0oC
Temperature of the water after heating T2 35.0oC
Mass of empty burner M1 28.3g
Mass of empty burner + ethanol before igniting M2 29.1g
Mass of empty burner + ethanol after igniting M3 28.7g

 

Sample calculations:

1.Calculate:

(a) ∆T the change in temperature

T = T2T1   => (35.0oC25.0oC) = 10.0oC

 

(b) the mass of ethanol used in burning

mass of ethanol used = M2 M1 => 29.1g 28.7g = 0.4g

 

(c) the number of moles of ethanol used in burning

moles of ethanol =   mass used => 0.4  = 0.0087 /8.7 x 10-3 moles
molar mass of ethanol 46

 

  1. Given that the specific heat capacity of water is 4.2 kJ-1kg-1K-1,determine the heat produced during the burning.

Heat produced H = mass of water(m) x specific heat capacity (c)x T

 

=> 20 x 4.2 x 10 = 840 Joules = 0.84 kJ 1000

3.Calculate the molar heat of combustion of ethanol

Molar heat of combustion ∆Hc = Heat produced ∆H

Number of moles of fuel

 

=> 0.84 kJ = 96.5517 kJmole-1 0.0087 /8.7 x 10-3 moles

 

4.List two sources of error in the above experiment.

 

(i)Heat loss to the surrounding lowers the practical value of the molar heat of combustion of ethanol.

 

13

 

A draught shield tries to minimize the loss by protecting wind from wobbling the flame.

(ii) Heat gain by reaction vessels/beaker lowers T and hence Hc

 

5.Calculate the heating value of  the fuel.

Heating value = molar heat of combustion => 96.5517 kJmole-1 = 2.0989 kJg-1

Molar mass of fuel 46 g

 

6.Explain other factors used to determine the choice of fuel for domestic and industrial use.

 

  • availability and affordability-some fuels are more available cheaply in rural than in urban areas at a lower cost.

 

(ii)cost of storage and transmission-a fuel should be easy to transport and store safely. e.g LPG is very convenient to store and use. Charcoal and wood are bulky.

 

  • environmental effects Most fuels after burning produce carbon(IV) oxide gas as a byproduct. Carbon(IV) oxide gas is green house gas that causes global warming. Some other fuel produce acidic gases like sulphur(IV) oxide ,and nitrogen(IV) oxide. These gases cause acid rain. Internal combustion engines exhaust produce lead vapour from leaded petrol and diesel. Lead is carcinogenic.

 

  • ignition point-The temperature at which a fuel must be heated before it burns in air is the ignition point. Fuels like petrol have very low ignition point, making it highly flammable. Charcoal and wood have very high ignition point.

 

7.Explain the methods used to reduce pollution from common fuels. (i)Planting trees-Plants absorb excess carbon(IV)oxide for photosynthesis

 

and release oxygen gas to the atmosphere.

 

(ii)using catalytic converters in internal combustion engines that convert harmful/toxic/poisonous gases like carbon(II)oxide and nitrogen(IV)oxide to harmless non-poisonous carbon(IV)oxide, water and nitrogen gas by using platinum-rhodium catalyst along the engine exhaust pipes.

 

Further practice calculations

 

1.Calculate the heating value of methanol CH3OH given that 0.87g of the fuel burn in air to raise the temperature of 500g of water from 20oC to 27oC.(C-12.0,H=1.0 O=16.0).

Moles of methanol used = Mass of methanol used => 0.87 g = 0.02718 moles

Molar mass of methanol 32

Heat produced H = mass of water(m) x specific heat capacity (c)x T => 500 x 4.2 x 7 = 14700 Joules = 14.7 kJ

1000

 

14

 

Molar heat of combustion ∆Hc = Heat produced ∆H
Number of moles of fuel
=> 14.7 kJ = 540.8389 kJmole-1

 

0.02718 moles

Heating value = molar heat of combustion => 540.8389 kJmole-1 = 16.9012 kJg-1

Molar mass of fuel 32 g

 

  1. 0 g of carbon burn in excess air to raise the temperature of 400g of water by 18oC.Determine the molar heat of combustion and hence the heating value of carbon(C-12.0,).
Moles of carbon used = Mass of carbon used => 1.0 g = 0.0833 moles
Molar mass of carbon 12

Heat produced H = mass of water(m) x specific heat capacity (c)x T => 400 x 4.2 x 18 = 30240 Joules = 30.24 kJ

1000

Molar heat of combustion ∆Hc = Heat produced ∆H
Number of moles of fuel
=> 30.24 kJ = 363.0252 kJmole-1

 

0.0833 moles

Heating value = molar heat of combustion => 363.0252 kJmole-1= 30.2521 kJg-1

Molar mass of fuel 12 g

 

(c)Standard enthalpy/heat of displacement ∆Hᶿd

 

The molar standard enthalpy/heat of displacement ∆Hᶿd is defined as the energy/heat change when one mole of a substance is displaced from its solution. A displacement reaction takes place when a more reactive element/with less electrode potential Eᶿ / negative Eᶿ /higher in the reactivity/electrochemical series remove/displace another with less reactive element/with higher electrode potential Eᶿ / positive Eᶿ /lower in the reactivity/electrochemical series from its solution.e.g.

 

(i)Zn(s) + CuSO4(aq) -> Cu(s) + ZnSO4(aq)

Ionically: Zn(s) + Cu2+(aq) -> Cu(s) + Zn2+ (aq)

(ii)Fe(s) + CuSO4(aq) -> Cu(s) + FeSO4(aq)

Ionically: Fe(s) + Cu2+(aq) -> Cu(s) + Fe2+ (aq)

(iii)Pb(s) + CuSO4(aq) -> Cu(s) + PbSO4(s)

 

This reaction stops after some time as insoluble PbSO4(s) coat/cover unreacted lead.

 

(iv)Cl2(g) + 2NaBr(aq) -> Br2(aq) + 2NaCl(aq)

Ionically: Cl2(g)+ 2Br (aq) -> Br2(aq) + 2Cl (aq)

 

15

 

Practically, a displacement reaction takes place when a known amount /volume of a solution is added excess of a more reactive metal.

 

To determine the molar standard enthalpy/heat of displacement(∆Hᶿd) of copper

 

Procedure

 

Place 20cm3 of 0.2M copper(II)sulphate(VI)solution into a 50cm3 plastic beaker/calorimeter. Determine and record the temperature of the solution T1.Put all the Zinc powder provided into the plastic beaker. Stir the mixture using the thermometer. Determine and record the highest temperature change to the nearest 0.5oC- T2 . Repeat the experiment to complete table 1 below

Table 1

 

Experiment I II
Final temperature of solution(T2) 30.0oC 31.0oC
Final temperature of solution(T1) 25.0oC 24.0oC
Change in temperature(T) 5.0 6.0

 

Questions

1.(a) Calculate:

 

(i)average ∆T

AverageT = change in temperature in experiment I and II =>5.0 + 6.0 = 5.5oC

2

 

(ii)the number of moles of solution used

Moles used = molarity x volume of solution   =    0.2 x 20    = 0.004 moles

 

1000                                              1000

 

(iii)the enthalpy change H for the reaction

Heat produced H = mass of solution(m) x specific heat capacity (c)x T => 20 x 4.2 x 5.5 = 462 Joules = 0.462 kJ

1000

 

(iv)State two assumptions made in the above calculations.

Density of solution = density of water = 1gcm-3

 

Specific heat capacity of solution=Specific heat capacity of solution=4.2 kJ-1kg-1K This is because the solution is assumed to be infinite dilute.

 

  1. Calculate the enthalpy change for one mole of displacement of Cu2+ (aq) ions.

Molar heat of displacement ∆Hd = Heat produced ∆H

 

16

 

Number of moles of fuel

=>             0.462 kJ              = 115.5 kJmole-1

0.004

 

3.Write an ionic equation for the reaction taking place.

Zn(s) + Cu2+(aq) ->    Cu(s) + Zn2+(aq)

 

4.State the observation made during the reaction.

 

Blue colour of copper(II)sulphate(VI) fades/becomes less blue/colourless.

Brown solid deposits are formed at the bottom of reaction vessel/ beaker.

 

5.Illustrate the above reaction using an energy level diagram.

 

Zn(s) + Cu2+(aq)
Energy H = -115.5 kJmole-1
(kJ)
Cu(s) + Zn2+(aq)
Reaction progress/path/coordinates

 

 

 

  1. Iron is less reactive than Zinc. Explain the effect of using iron instead of

Zinc on the standard molar heat of displacement ∆Hd of copper(II)sulphate (VI) solution.

 

No effect. Cu2+ (aq) are displaced from their solution.The element used to displace it does not matter.The reaction however faster if a more reactive metal is used.

 

7.(a)If the standard molar heat of displacement ∆Hd of copper(II)sulphate (VI) solution is 209kJmole-1 calculate the temperature change if 50cm3 of

0.2M solution was displaced by excess magnesium.
Moles used = molarity x volume of solution =  0.2 x 50 = 0.01 moles
1000 1000
Heat produced ∆H = Molar heat of displacement ∆Hd x Number of moles
=>209kJmole-1x 0.01 moles = 2.09 kJ
T (change in temperature) = Heat produced ∆H
Molar heat of displacement ∆Hd x Number of moles
=>2.09 kJ = 9.9524Kelvin
0.01 moles

 

(b)Draw an energy level diagram to show the above energy changes

 

 

17

 

Mg(s) + Cu2+(aq)
Energy H = -209 kJmole-1
(kJ)

Cu(s) + Mg2+(aq)

 

Reaction progress/path/coordinates

 

  1. The enthalpy of displacement ∆Hd of copper(II)sulphate (VI) solution is 12k6kJmole-1.Calculate the molarity of the solution given that 40cm3 of this solution produces 2.204kJ of energy during a displacement reaction with

 

excess iron filings. Heat produced ∆H
Number of moles =
Molar heat of displacement ∆Hd
=>2.204 kJ = 0.0206moles
126 moles
Molarity of the solution  = moles x 1000
Volume of solution used
= 0.0206moles x 1000 = 0.5167 M
40

 

  1. If the molar heat of displacement of Zinc(II)nitrate(V)by magnesium powder is 25.05kJmole-1 ,calculate the volume of solution which must be added 0.5 moles solution if there was a 3.0K rise in temperature.

 

Heat produced ∆H = Molar heat of displacement ∆Hd x Number of moles

=>25.08kJmole-1x 0.5 moles = 1.254 kJ x 1000 =1254J
Mass of solution (m)  = Heat produced H
specific heat capacity (c)x T
=> 1254J = 99.5238 g
4.2 x 3
Volume = mass x density = 99.5238 g x 1 = 99.5238cm3

 

Note: The solution assumes to be too dilute /infinite dilute such that the density and specific heat capacity is assumed to be that of water.

 

Graphical determination of the molar enthalpy of displacement of copper

 

Procedure:

 

 

18

 

Place 20cm3 of 0.2M copper(II)sulphate (VI) solution into a calorimeter/50cm3 of plastic beaker wrapped in cotton wool/tissue paper. Record its temperature at time T= 0.

 

Stir the solution with the thermometer carefully and continue recording the

temperature after every 30 seconds .

Place all the (1.5g) Zinc powder provided.

 

Stir the solution with the thermometer carefully and continue recording the temperature after every 30 seconds for five minutes.

Determine the highest temperature change to the nearest 0.5oC.

 

Sample results

Time oC 0.0 30.0 60.0 90.0 120.0 150.0 180.0 210.0 240.0 270.0
Temperature 25.0 25.0 25.0 25.0 25.0 xxx 36.0 35.5 35.0 34.5
Sketch graph of temperature against time
36.5
Extrapolation T
Temperature point
oC
130 Time(seconds)

Questions

  1. Show and determine the change in temperature ∆T

From a well constructed graph T= T2 T1 at 150 second by extrapolation T = 36.525.0 = 11.5oC

 

2.Calculate the number of moles of copper(II) sulphate(VI)used given the molar heat of displacement of Cu2+ (aq)ions is 125kJmole-1

 

Heat produced H = mass of solution(m) x specific heat capacity (c)x T => 20 x 4.2 x 11.5 = 966 Joules = 0.966 kJ

1000

 

Number of moles =              Heat produced ∆H Molar heat of displacement ∆Hd

 

 

=>.966 kJ = 0.007728moles
7.728 x 10-3moles
125 moles

 

  1. What was the concentration of copper(II)sulphate(VI) in moles per litre.

 

 

19

 

Molarity =  moles x 1000 => 7.728 x 10-3moles x 1000 = 0.3864M
Volume used 20

 

4.The actual concentration of copper(II)sulphate(VI) solution was

0.4M.Explain the differences between the two.

 

Practical value is lower than theoretical. Heat/energy loss to the surrounding and that absorbed by the reaction vessel decreases T hence lowering the practical number of moles and molarity against the theoretical value

 

 

 

5.a) In an experiment to determine the molar heat of reaction when magnesium displaces copper ,0.15g of magnesium powder were added to 25.0cm3 of 2.0M copper (II) chloride solution. The temperature of copper (II) chloride solution was 25oC.While that of the mixture was 43oC.

 

i)Other than increase in temperature, state and explain the observations which were made during the reaction.(3mks)

 

ii)Calculate the heat change during the reaction (specific heat capacity of the solution = 4.2jg-1k-1and the density of the solution = 1g/cm3(2mks)

 

iii)Determine the molar heat of displacement of copper by magnesium.(Mg=24.0). iv)Write the ionic equation for the reaction.(1mk)

 

v)Sketch an energy level diagram for the reaction.(2mks)

 

 

 

b)Use the reduction potentials given below to explain why a solution containing copper ions should not be stored in a container made of zinc.

 

 

Zn2+(aq) + 2e Cu2+(aq) + 2e

 

 

->

->

 

 

Zn(s); Cu(s);

 

 

Eø = -0.76v Eø = +0.34v

 

 

 

 

(2mks)

 

 

 

 

 

(c)Standard enthalpy/heat of neutralization ∆Hᶿn

 

The molar standard enthalpy/heat of neutralization ∆Hᶿn is defined as the energy/heat change when one mole of a H+ (H3O+)ions react completely with one mole of OH ions to form one mole of H2O/water.

 

Neutralization is thus a reaction of an acid /H+ (H3O+)ions with a base/alkali/ OHions to form salt and water only.

 

 

20

 

Strong acids/bases/alkalis are completely dissociated to many free ions(H+ /H3O+ and OH ions).

 

Weak acids/bases/alkalis are partially dissociated to few free ions(H+ (H3O+ and OH ions) and exist more as molecules.

Neutralization is an exothermic(-H) process.The enrgy produced during neutralization depend on the amount of free ions (H+ H3O+ and OH)ions existing in the acid/base/alkali reactant:

 

(i)for weak acid-base/alkali neutralization,some of the energy is used to dissociate /ionize the molecule into free H+ H3O+ and OH ions therefore the overall energy evolved is comparatively lower/lesser/smaller than strong acid / base/ alkali neutralizations.

 

  • (i)for strong acid/base/alkali neutralization, no energy is used to dissociate /ionize since molecule is wholly/fully dissociated/ionized into free H+ H3O+ and OHThe overall energy evolved is comparatively higher/more than weak acid-base/ alkali neutralizations. For strong acid-base/alkali neutralization, the enthalpy of neutralization is constant at about 57.3kJmole-1 irrespective of the acid-base used. This is because ionically:

 

OH(aq)+  H+(aq)  ->   H2O(l)  for any wholly dissociated acid/base/alkali

 

 

 

To determine the molar enthalpy of neutralization ∆Hn of Hydrochloric acid Procedure

 

 

Place 50cm3 of 2M hydrochloric acid into a calorimeter/200cm3 plastic beaker wrapped in cotton wool/tissue paper. Record its temperature T1.Using a clean measuring cylinder, measure another 50cm3 of 2M sodium hydroxide. Rinse the bulb of the thermometer in distilled water. Determine the temperature of the sodium hydroxide T2.Average T2 andT1 to get the initial temperature of the mixture T3.

 

Carefully add all the alkali into the calorimeter/200cm3 plastic beaker wrapped in cotton wool/tissue paper containing the acid. Stir vigorously the mixture with the thermometer.

 

Determine the highest temperature change to the nearest 0.5oC T4 as the final temperature of the mixture. Repeat the experiment to complete table 1.

 

Table I . Sample results

 

Experiment I II
Temperature of acid T1 (oC) 22.5 22.5
Temperature of base T2 (oC) 22.0 23.0
Final temperature of solution T4(oC) 35.5 36.0
Initial temperature of solution T3(oC) 22.25 22.75
21

 

 

Temperature change( T5) 13.25 13.75

 

(a)Calculate

T6 the average temperature change

T6 =    13.25 +13.75    = 13.5 oC

2

(b)Why should the apparatus be very clean?

 

Impurities present in the apparatus reacts with acid /base lowering the overall temperature change and hence Hᶿn.

 

(c)Calculate the:
(i)number of moles of the acid used
number of moles = molarity x volume => 2 x 50   = 0.1moles
1000 1000

 

(ii)enthalpy change H of neutralization.

H = (m)mass of solution(acid+base) x (c)specific heat capacity of solution x

T(T6) => (50 +50) x 4.2 x 13.5 = 5670Joules = 5.67kJ
(iii) the molar heat of neutralization the acid.
Hn = Enthalpy change ∆H => 5.67kJ = 56.7kJ mole-1
Number of moles 0.1moles

 

(c)Write the ionic equation for the reaction that takes place

OH(aq)+ H+(aq) ->      H2O(l)

 

(d)The theoretical enthalpy change is 57.4kJ. Explain the difference with the results above.

The theoretical value is higher

Heat/energy loss to the surrounding/environment lowers T/T6 and thus Hn Heat/energy is absorbed by the reaction vessel/calorimeter/plastic cup

lowers T and hence Hn

 

(e)Compare the Hn of the experiment above with similar experiment repeated with neutralization of a solution of:

 

  • potassium hydroxide with nitric(V) acid

The results would be the same/similar.

 

Both are neutralization reactions of strong acids and bases/alkalis that are fully /wholly dissociated into many free H+ / H3O+ and OH ions.

 

  • ammonia with ethanoic acid

 

22

 

The results would be lower/∆Hn would be less.

 

Both are neutralization reactions of weak acids and bases/alkalis that are partially /partly dissociated into few free H+ / H3O+ and OH ions. Some energy is used to ionize the molecule.

 

(f)Draw an energy level diagram to illustrate the energy changes H2 H+ (aq)+OH (aq)

Energy

(kJ)

 

H = -56.7kJ

 

H1                                                                 H2O (l)

 

Reaction path/coordinate/progress

 

Theoretical examples

 

1.The molar enthalpy of neutralization was experimentary shown to be 51.5kJ per mole of 0.5M hydrochloric acid and 0.5M sodium hydroxide. If the volume of sodium hydroxide was 20cm3, what was the volume of hydrochloric acid used if the reaction produced a 5.0oC rise in temperature?

 

Working:
Moles of sodium hydroxide = molarity x volume => 0.5 M x 20cm3 = 0.01 moles
1000 1000
Enthalpy change H = H n => 51.5 = 0.515kJ
Moles sodium hydroxide 0.01 moles

 

Mass of base + acid =           Enthalpy change H in Joules Specific heat capacity x T

=>                      0.515kJ    x 1000         =    24.5238g

 

4.2 x 5

 

Mass/volume of HCl = Total volume volume of NaOH =>24.5238 – 20.0 = 4.5238 cm3

 

  1. Hn of potassium hydroxide was practically determined to be

56.7kJmole-1.Calculate the molarity of 50.0 cm3 potassium hydroxide used to neutralize 25.0cm3 of dilute sulphuric(VI) acid raising the temperature of the solution from 10.0oC to 16.5oC.

 

H = (m)mass of solution(acid+base) x (c)specific heat capacity of solution xT

 

23

 

=> (50 +25) x 4.2 x 6.5 = 2047.5Joules
Moles potassium hydroxide =Enthalpy change H
Hn
2047.5Joules    = 0.0361 moles
56700Joules
Molarity of KOH = moles x 1000 => 0.0361 moles x 1000  = 0.722M
Volume used 50cm3

 

3.Determine the specific heat capacity of a solution of a solution mixture of 50.0cm3 of 2M potassium hydroxide neutralizing 50.0cm3 of 2M nitric(V) acid if a 13.25oC rise in temperature is recorded.(1mole of potassium hydroxide produce 55.4kJ of energy)

 

Moles of potassium hydroxide = molarity KOH x volume 1000

=> 2 M x 50cm3 = 0.1 moles

1000
Enthalpy change H = Hn    x Moles potassium hydroxide
=> 55.4kJ x 0.1 moles = 5.54kJ x 1000=5540Joules
Specific heat capacity = Enthalpy change H in Joules
Mass of base + acid x T
=> 5540 4.1811J-1g-1K-1

 

(50+50) x 13.25

 

Graphically Hn can be determined as in the example below:

 

Procedure

 

Place 8 test tubes in a test tube rack .Put 5cm3 of 2M sodium hydroxide solution into each test tube.

 

Measure 25cm3 of 1M hydrochloric acid into 100cm3 plastic beaker.

 

Record its initial temperature at volume of base =0. Put one portion of the base into the beaker containing the acid.

 

Stir carefully with the thermometer and record the highest temperature change to the nearest 0.5oC.

 

Repeat the procedure above with other portions of the base to complete table 1 below

 

Table 1:Sample results.

 

24

 

Volume of acid(cm3) 25.0 25.0 25.0 25.0 25.0 25.0 25.0 25.0 25.0
Volume of alkali(cm3) 0 5.0 10.0 15.0 20.0 25.0 30.0 35.0 40.0
Final temperature(oC) 22.0 24.0 26.0 28.0 28.0 27.0 26.0 25.0 24.0
Initial temperature(oC) 22.0 22.0 22.0 22.0 22.0 22.0 22.0 22.0 22.0
Change in temperature 0.0 2.0 4.0 6.0 6.0 5.0 4.0 3.0 2.0

 

(a)Complete the table to determine the change in temperature.

 

(b)Plot a graph of volume of sodium hydroxide against temperature change.

 

 

 

28.7=T2  temperature(oC)

 

T

 

 

 

  1. oC T1

 

Volume of sodium hydroxide(cm3)

 

From the graph show and determine :

(i)the highest temperature change ∆T

T =T2-T1  => highest temperature-T2 (from extrapolating a correctly plotted

 

graph) less lowest temperature at volume of base=0 :T1 =>T = 6.7 0.0 = 6.70C

 

(ii)the volume of sodium hydroxide used for complete neutralization From a correctly plotted graph 16.75cm3

 

(c)Calculate the number of moles of the alkali used

Moles NaOH = molarity x volume =>2M x 16.75cm3  = 0.0335 moles
1000 1000

 

(d)Calculate H for the reaction

H = mass of solution(acid+base) x c xT

 

=>(25.0 + 16.75) x 4.2 x 6.7 = 1174.845 J = 1.174845kJ 1000

 

 

(e)Calculate the molar enthalpy of neutralization of the alkali.

∆Hn  = ∆H n = 1.174845kJ = 35.0701kJ
Number of moles 0.0335
25

 

(d)Standard enthalpy/heat of solution ∆Hᶿs

 

The standard enthalpy of solution Hᶿsis defined as the energy change when one mole of a substance is dissolve in excess distilled water to form an infinite dilute solution. An infinite dilute solution is one which is too dilute to be diluted further.

 

Dissolving a solid involves two processes:

 

  • breaking the crystal of the solid into free ions(cations and anion).This process is the opposite of the formation of the crystal itself. The energy required

 

to form one mole of a crystal structure from its gaseous ions is called Lattice energy/heat/enthalpy of lattice (Hl). Lattice energy /heat/enthalpy of lattice (Hl) is an endothermic process (+∆Hl).

 

The table below shows some Hl in kJ for the process MX(s) -> M+ (g) + X (g)
Li Na K Ca Mg
F +1022 +900 +800 +760 +631
Cl +846 +771 +690 +2237 +2493
Br +800 +733 +670 +2173 +2226

 

(ii)surrounding the free ions by polar water molecules. This process is called hydration. The energy produced when one mole of ions are completely hydrated is called hydration energy/ heat/enthalpy of hydration(Hh).Hydration energy /enthalpy of hydration(Hh) is an exothermic process(Hh).

 

The table below shows some Hh in kJ for some ions;

ion Li+ Na+ K+ Mg2+ Ca2+ F Cl Br
Hh -1091 -406 -322 -1920 -1650 -506 -364 -335

 

The sum of the lattice energy +Hl (endothermic) and hydration energy Hh (exothermic) gives the heat of solutionHs

Hs =Hl +Hh

Note

Since Hl is an endothermic process and ∆Hh is an exothermic process then ∆Hs is:

(i)exothermic if Hl is less than Hh and hence a solid dissolve easily in

water.

(ii)endothermic if Hl is more than Hh and hence a solid does not dissolve easily in water.

 

(a)Dissolving sodium chloride crystal/s:

 

(i) NaCl –—breaking the crystal into free ions–->  Na +(g)+ Cl(g)  Hl =+771 kJ

 

26

 

(ii) Hydrating the ions;

Na +(g) + aq -> Na(aq) Hh = – 406 kJ Cl(g) + aq -> Cl(aq) Hh = – 364 kJ

 

Hs =Hh +Hs -> (- 406 kJ + – 364 kJ) + +771 kJ = + 1.0 kJmole-1 NaCl does not dissolve easily in water because overall Hs is endothermic

 

Solubility of NaCl therefore increases with increase in temperature.

 

Increase in temperature increases the energy to break the crystal lattice of NaCl to free Na +(g)+ Cl(g)

 

(b)Dissolving magnesium chloride crystal/s//  MgCl2 (s) ->MgCl2 (aq)

 

  • MgCl2 -breaking the crystal into free ions-->Mg 2+(g)+ 2Cl(g) Hl =+2493 kJ
  • Hydrating the ions;

 

Mg 2+(g) + aq -> Mg 2+(g) (aq) ∆Hh = – 1920 kJ 2Cl(g) + aq -> 2Cl(aq) ∆Hh = (- 364 x 2) kJ

Hs =Hh +Hs -> (- 1920 kJ + (- 364 x 2 kJ)) + +2493 kJ = 155.0 kJmole-1

 

MgCl2 (s) dissolve easily in water because overall Hs is exothermic .

Solubility of MgCl2 (s) therefore decreases with increase in temperature.

 

(c)Dissolving Calcium floride crystal/s//  CaF2 (s) -> CaF2 (aq)

  • CaF2 –>Ca 2+(g)+ 2F(g) ∆Hl =+760 kJ
  • Hydrating the ions;

 

Ca 2+(g) + aq -> Ca 2+(g) (aq) ∆Hh = – 1650 kJ 2F(g) + aq -> 2F(aq) ∆Hh = (- 506 x 2) kJ

Hs =Hh +Hs -> (- 1650 kJ + (- 506 x 2 kJ)) + +760 kJ = 1902.0 kJmole-1

 

CaF2 (s) dissolve easily in water because overall Hs is exothermic .

Solubility of CaF2 (s) therefore decreases with increase in temperature.

 

(d)Dissolving magnesium bromide crystal/s//  MgBr2 (s) ->MgBr2 (aq)

 

  • MgCl2 -breaking the crystal into free ions-->Mg 2+(g)+ 2Br(g) ∆Hl =+2226 kJ
  • Hydrating the ions;

 

Mg 2+(g) + aq -> Mg 2+(g) (aq) ∆Hh = – 1920 kJ 2Br(g) + aq -> 2Br(aq) ∆Hh = (- 335x 2) kJ

Hs =Hh +Hs -> (- 1920 kJ + (- 335 x 2 kJ)) + +2226 kJ = 364.0 kJmole-1

 

MgBr2 (s) dissolve easily in water because overall Hs is exothermic .

Solubility of MgBr2(s) therefore decreases with increase in temperature.

 

27

 

Practically the heat of solution can be determined from dissolving known amount /mass/volume of solute in known mass /volume of water/solvent.

 

From the temperature of solvent before and after dissolving the change in temperature(T) during dissolution is determined.

 

To determine the ∆Hs ammonium nitrate

 

Place 100cm3 of distilled water into a plastic beaker/calorimeter. Determine its temperature and record it at time =0 in table I below.

 

Put all the 5.0g of ammonium nitrate (potassium nitrate/ammonium chloride can also be used)provided into the plastic beaker/calorimeter, stir using a thermometer and record the highest temperature change to the nearest 0.5oCafter every ½ minute to complete table I.

 

Continue stirring the mixture throughout the experiment.

Sample results: Table I

 

Time (minutes) 0.0 ½ 1 1 ½ 2 2 ½ 3 3 ½
Temperature()oC 22.0 21.0 20.0 19.0 19.0 19.5 20.0 20.5

 

(a)Plot a graph of temperature against time(x-axis)

 

 

 

 

 

 

 

22.0=T1

 

temperature(oC)

T

 

 

 

 

 

18.7. oC T1

 

 

Time (minutes)

 

 

(b)From the graph show and determine the highest temperature change ∆T

T =T2-T1 => lowest temperature-T2 (from extrapolating a correctly plotted graph) less highest temperature at volume of base=0 :T1

 

28

 

=>T =18.7 22.0 =   3.30C

 

(c)Calculate the number of moles of ammonium nitrate(V) used

Moles NH4NO3 = mass used => 5.0 = 0.0625 moles
Molar mass 80

 

(d)Calculate H for the reaction

H = mass of water x c xT

 

->100 x 4.2 x 3.3 = +1386 J = +1.386kJ 1000

 

 

(e)Calculate the molar enthalpy of dissolution of ammonium nitrate(V).

∆Hs  = ∆H = +1.386kJ = + 22.176kJ mole-1
Number of moles 0.0625 moles

 

(f)What would happen if the distilled water was heated before the experiment was performed.

 

The ammonium nitrate(V)would take less time to dissolves. Increase in temperature reduces lattice energy causing endothermic dissolution to be faster

 

(g)Illustrate the process above in an energy level diagram

 

NH4+ (g) + NO3(g)

 

+H

NH4+ (aq)+NO3(aq)

Energy(kJ)

 

+H                                          H = -22.176kJ

 

 

 

NH4NO3(s)

 

 

 

Reaction path /progress/coordinate

 

  • 100cm3 of distilled water at 25oC was added carefully 3cm3 concentrated sulphuric(VI)acid of density 1.84gcm-3.The temperature of the mixture rose from 250C to 38oCalculate the molar heat of solution of sulphuric(VI)acid (S=32.0,H=1.0,0=16.0)

 

Working

 

29

 

Molar mass of H2SO4 = 98g
Mass of H2SO4= Density x volume => 1.84gcm-3 x 3cm3 = 5.52 g
Mass of H2O  = Density x volume => 1.00gcm-3 x 100cm3 = 100 g
Moles of H2SO4= mass => 5.52 g =   0.0563 moles
Molar mass of H2SO4 98g

 

Enthalpy change H= (mass of acid + water) x specific heat capacity of water x T => (100 +5.52 g) x 4.2 x 13oC = 5761.392 J = 5.761392 kJ

1000

 

Hs of H2SO4= H => 5.761392 kJ = -102.33378kJmoles-1
Moles of H2SO4 0.0563 moles

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(e)Standard enthalpy/heat of formation  ∆Hᶿf

 

The molar enthalpy of formation ∆Hᶿf is defined as the energy change when one mole of a compound is formed from its elements at 298K(25oC) and 101325Pa(one atmosphere)pressure. Hᶿf is practically difficult to determine in a school laboratory.

 

It is determined normally determined by applying Hess law of constant heat summation.

 

Hess law of constant heat summation states that the total enthalpy/heat/energy change of a reaction is the same regardless of the route taken from reactants to products at the same temperature and pressure.

 

Hess law of constant heat summation is as a result of a series of experiments done by the German Scientist Henri Hess(1802-1850).

 

He found that the total energy change from the reactants to products was the same irrespective of the intermediate products between. i.e.

 

A(s)    —H1–>C(s)    =       A(s) —H2–>B(s)–H3–>C(s)

 

30

 

A(s)   ∆H2                                      B(s)

 

H1                                         H3

 

 

C(s)

 

The above is called an energy cycle diagram. It can be used to calculate any of the missing energy changes since:

  • ∆H1 =∆H2   + ∆H3
  • ∆H2 =∆H1 +  –∆H3

 

  • ∆H3 = – ∆H1 +  ∆H2

 

Examples of applying Hess law of constant heat summation

 

1.Calculate the molar enthalpy of formation of methane (CH4) given that Hᶿc of carbon-graphite is -393.5kJmole-1,Hydrogen is -285.7 kJmole-1 and that of methane is 890 kJmole-1

 

Working

Carbon-graphite ,hydrogen and oxygen can react to first form methane.

Methane will then burn in the oxygen present to form carbon(IV)oxide and water.

Carbon-graphite can burn in the oxygen to form carbon(IV)oxide.

Hydrogen can burn in the oxygen to form water.

C(s)+ 2H2 (g)+2O2 (g) —H1–> CH4(g) +2O2(g) —H2–> CO2(g)+2H2O(l)

C(s)+ 2H2 (g)+2O2 (g) —H3–> CO2(g)+2H2O(l)

 

Energy cycle diagram

C(s) +    2H2 (g) + 2O2(g)              H1=Hᶿc =-890.4kJCH4(g)+2O2(g)

 

 

 

 

H3=Hᶿc =-393.5kJ   ∆H3=Hᶿc =-285.7kJ x 2      H2= ∆Hᶿf= x

 

 

CO2(g)          +       2H2O(l)

 

Substituting:

H3   =H1  + ∆H2

-393.5 + (-285.7 x 2) = -890.4kJ + x

x = -74.5 kJ

 

31

 

Heat of formation ∆Hᶿf  CH4 = -74.5 kJmole-1

 

  1. Calculate the molar enthalpy of formation of ethyne (C2H2) given : Hᶿc of carbon-graphite = -394kJmole-1,Hydrogen = -286 kJmole-1 , (C2H2) = -1300 kJmole-1

 

Working

Carbon-graphite ,hydrogen and oxygen can react to first form ethyne.

 

Ethyne will then burn in the oxygen present to form carbon(IV)oxide and water.

Carbon-graphite can burn in the oxygen to form carbon(IV)oxide.

 

Hydrogen can burn in the oxygen to form water.
2C(s)+ H2 (g)+2 ½ O2 (g) —H1–> C2 H2 (g) +2 ½ O2(g) —H2–> CO2(g)+H2O(l)
2C(s)+ H2 (g)+ 2 ½ O2 (g) —H3–> 2CO2(g)+H2O(l)
Energy cycle diagram
2C(s) + H2 (g) +2½O2(g) H1=Hᶿf =x C2 H2+2½O2(g)

 

 

 

 

H3=∆Hᶿc =-394kJx 2   ∆H3=∆Hᶿc =-286kJ       H2= ∆Hᶿc= -1300kJ

 

 

2CO2(g)          +         H2O(l)

 

 

 

Substituting:

H3   = ∆H1  + ∆H2

 

  • -394 x 2) + -286 = -1300kJ + x x = +244 kJ

Heat of formation ∆Hᶿf  CH4 = +244 kJmole-1

 

  1. Calculate the molar enthalpy of formation of carbon(II)oxide (CO) given : Hᶿc of carbon-graphite = -393.5kJmole-1,Hᶿc of carbon(II)oxide (CO)= -283 kJmole-1

 

Working

Carbon-graphite reacts with oxygen first to form carbon (II)oxide (CO).

Carbon(II)oxide (CO) then burn in the excess oxygen to form carbon(IV)oxide.

 

Carbon-graphite can burn in excess oxygen to form carbon (IV) oxide.

 

 

32

 

C(s)+  ½O2 (g) H1 > CO (g) + ½ O2(g)  H2—- > CO2(g)
C(s)+ O2 (g) — H3–> CO2(g)
Energy cycle diagram
C(s)  + ½O2(g) H1=∆Hᶿf =x CO+½O2(g)

 

 

 

 

H3=∆Hᶿc =-393.5kJ                                               H2= ∆Hᶿc= -283kJ

 

 

CO2(g)

 

Substituting:

 

H3   = ∆H1  + ∆H2

-393.5kJ = -283kJ + x

x = -110 kJ

Heat of formation ∆Hᶿf  CO = -110 kJmole-1

 

4.Study the information below:

H2(g) + ½ O2(g) -> H2O(l) ∆H1=-286 kJmole-1 C(s) + O2(g) -> CO2(g) ∆H2=-393 kJmole-1

2C(s) + H2(g) + ½ O2(g)                                   ->C2H5OH(l)    ∆H3=-277 kJmole-1

Use the information to calculate the molar enthalpy of combustion H4 of ethanol

 

Energy cycle diagram

 

2C(s) + 3H2 (g) +3½O2(g)          H3=Hᶿf =-227kJ  C2 H5OH +3O2(g)

 

 

 

 

H2=∆Hᶿc =-394kJx 2   ∆H1=∆Hᶿc =-286kJx 3       H4= ∆Hᶿc= x

 

 

2CO2(g)         +         3H2O(l)

 

Substituting:

H1   +   ∆H2  = ∆H3 +H4

 

  • -394 x 2) + -286 x 3 = -277 + x H4 = -1369 kJ

 

33

 

Heat of combustion ∆Hᶿc  C2H5OH = -1369 kJmole-1

 

5.Given the following information below:

CuSO4(s) + (aq)                                      -> CuSO4(aq)   ∆H=-66.1 kJmole-1

 

CuSO4(s) + (aq) + 5H2O(l)-> CuSO4 .5H2O (aq) ∆H=-77.4 kJmole-1 Calculate ∆H for the reaction;

 

CuSO4(aq) + 5H2O-> CuSO4 .5H2O (aq) ∆H=-77.4 kJmole-1 Working

 

CuSO4(s) + (aq)  + 5H2O(l)->  CuSO4(aq)+ 5H2O(l)-> CuSO4 .5H2O (aq)

CuSO4(s) + (aq)  + 5H2O(l)-> CuSO4 .5H2O (aq)

 

Energy cycle diagram

 

CuSO4(s) + (aq) + 5H2O(l)                H1=+66.1kJ   CuSO4(aq)+ 5H2O(l)

 

 

 

 

H3= =-77.4kJ                                               H2= x

 

 

CuSO4 .5H2O (aq)

 

Substituting:

 

H3   = ∆H2 +H1

  • -77.4kJ = x + +66.1kJ

 

∆H4 = -10.9 kJ

Heat of dissolution of CuSO4 = -10.9kJmole-1

 

Practically, Hess law can be applied practically as in the following examples

 

a)Practical example 1

Determination of the enthalpy of formation of CuSO4.5H2O

Experiment I

 

Weigh accurately 12.5 g of copper(II)sulphate(VI)pentahydrate. Measure 100cm3 of distilled water into a beaker. Determine its temperature T1 .Put all the crystals of the copper(II)sulphate(VI)pentahydrate carefully into the beaker. Stir using a thermometer and determine the highest temperature change T2 Repeat the procedure again to complete table 1.

 

Table 1:Sample results

 

 

34

 

Experiment I II
Highest /lowest temperature T2 27.0 29.0
Initial temperature T1 24.0 25.0
Change in temperature T 3.0 4.0

 

Experiment II

 

Weigh accurately 8.0g of anhydrous copper(II)sulphate(VI). Measure 100cm3 of distilled water into a beaker. Determine its temperature T1 .Put all the crystals of the copper(II)sulphate(VI)pentahydrate carefully into the beaker. Stir using a thermometer and determine the highest temperature change T2 Repeat the procedure again to complete table II.

 

Table II :Sample results

 

Experiment I II
Highest /lowest temperature T2 26.0 27.0
Initial temperature T1 25.0 25.0
Change in temperature T 1.0 2.0

 

Questions

(a)Calculate the average T in

(i)Table I

T= T2 -T1 => 3.0 +4.0  = 3.5 oC

2

(ii)Table II

T= T2 -T1 => 1.0 +2.0  = 1.5 oC

2

 

(b)Calculate the number of moles of solid used in:

(i)Experiment I

Moles of CuSO4.5H2O = Mass => 12.5 = 0.05 moles
Molar mass 250
(ii)Experiment II
Moles of CuSO4 = Mass => 8.0 = 0.05 moles
Molar mass 160

 

(c)Calculate the enthalpy change for the reaction in:

(i)Experiment I

Enthalpy change of CuSO4.5H2O= mass of Water(m) x c x T

=>100cm3 x 4.2 x 3.5 oC = -1.47kJ

 

1000

(ii)Experiment II

Enthalpy change of CuSO4 = mass of water(m) x c x T

35

 

=>100cm3 x 4.2 x 1.5 oC = -0.63kJ

 

 

 

(c)Calculate the molar enthalpy of solution CuSO4 .5H2O (s) form the results in (i)experiment I.

Hs = CuSO4.5H2O= H => -1.47kJ = 29.4kJ
Number of Moles 0.05 moles
(ii)experiment II.
Hs = CuSO4= H => -0.63kJ 12.6kJ
Number of Moles 0.05 moles

 

 

 

  • Using an energy level diagram, calculate the molar enthalpy change for the

 

reaction:

CuSO4 .5H2O (s) -> CuSO4(s) + 5H2O(l)

 

Energy cycle diagram

 

 

CuSO4(s) + (aq) + 5H2O(l)

 

 

H1=x

 

 

CuSO4. 5H2O (s)+ (aq)

 

 

 

 

 

H3= =-29.4kJ                                               H2= -12.6kJ

 

 

CuSO4 .5H2O (aq)

 

H3 =H1 +H2

=>-29.4kJ = -12.6kJ + x

 

=>-29.4kJ – (+12.6kJ) = x

x = 16.8kJ

 

 

 

b)Practical example II

Determination of enthalpy of solution of ammonium chloride

 

Theoretical information.

 

Ammonium chloride dissolves in water to form ammonium chloride solution. Aqueous ammonia can react with excess dilute hydrochloric acid to form ammonium chloride solution. The heat change taking place can be calculated from the heat of reactions:

 

36

 

  • NH3(aq) + HCl(aq) -> NH4Cl(s)
  • NH4Cl(s) + (aq) -> NH4Cl(aq)

 

  • NH3(aq) + HCl(aq) -> NH4Cl(aq)

 

Experiment procedure I

Measure 50cm3 of water into a 100cm3 beaker. Record its temperature T1 as initial temperature to the nearest 0.5oC in table I. Add exactly 5.0g of ammonium chloride crystals weighed carefully into the water. Stir and record the highest temperature change T2 as the final temperature change. Repeat the above procedure to complete table I.

 

Sample results TableI

 

Experiment I II
final temperature(oC) 19.0 20.0
initial temperature(oC) 22.0 22.0
temperature change T(oC) 3.0 2.0

 

Experiment procedure II

 

Measure 25cm3 of 2M aqueous ammonia into a 100cm3 beaker. Record its temperature T1 as initial temperature to the nearest 0.5oC in table II. Measure 25cm3 of 2M hydrochloric acid solution. Add the acid into the beaker containing aqueous ammonia. Stir and record the highest temperature change T2 as the final temperature change. Repeat the above procedure to complete table II.

 

Sample results:Table II

 

Experiment I II
final temperature(oC) 29.0 29.0
initial temperature(oC) 22.0 22.0
temperature change T(oC) 7.0 7.0

 

Sample Calculations:

(a)Calculate the average ∆T in

(i)Table I

T= T2 -T1 => –3.0 +-2.0   = 2.5 oC

2

(ii)Table II

T= T2 -T1 => 7.0 +7.0  = 7.0 oC

2

 

 

 

(b)Calculate the enthalpy change for the reaction in:

 

37

 

(i)Experiment I

Enthalpy change H = mass of Water(m) x c x T

 

=>50cm3 x 4.2 x 2.5 oC = +0.525kJ 1000

 

(ii)Experiment II

Enthalpy change of CuSO4 = mass of water(m) x c x T

=>25+25cm3 x 4.2 x 7 oC = +1.47kJ

1000

 

(c)Write the equation for the reaction taking place in:

(i)Experiment I

NH4Cl(s)  + (aq) -> NH4Cl(aq)

 

(ii)Experiment I

NH3(aq)   + HCl(aq)      -> NH4Cl(aq)

 

(d)Calculate the enthalpy change ∆H for the reaction:

NH3(g)                                                                                 + HCl(g) -> NH4Cl(s) given that:

 

 

(i) NH3(g)  + (aq) -> NH3(aq)

(ii) (aq)   + HCl(g) -> HCl(aq)

 

 

H= -40.3kJ

 

H= -16.45kJ

 

 

 

(e)Applying Hess Law of constant heat summation:

 

Energy level diagram

 

N2(g)                +    1½ H2(g)             + ½ Cl2            Hf    NH4Cl(s) + aq

 

+0.525kJ=H4

 

 

(aq) (aq)
– 40.3kJ=H1 -16.43kJ=H2
NH3 (aq)  + HCl(aq) -1.47kJ=H3 NH4Cl(s)
H1 +H2 +    H3 =   H4 + Hf
– 40.3kJ  + -16.43kJ  +  -1.47kJ   =  +0.525kJ + Hf
=>Hf = -58.865kJ.

 

Practice theoretical examples:

 

38

 

  1. Using an energy level diagram calculate the ∆Hs of ammonium chloride crystals given that.

Hf of NH3 (aq) = -80.54kJ mole-1Hf of HCl (aq) = -164.46kJ mole-1Hf of NH4Cl (aq) = -261.7483kJ mole-1Hs of NH4Cl (aq) = -16.8517kJ mole-1

 

N2(g) +    1½ H2(g)   + ½ Cl2             Hf=-261.7483kJ                    NH4Cl(s) + aq

 

 

 

 

x=Hs
(aq) (aq)
– 80.54kJ=H1 -164.46kJ=H2
NH3 (aq)  + HCl(aq) 16.8517kJ=H3 NH4Cl(s)
H1 +H2 +H3 =   H4 + Hf
  • 54kJ + -164.46kJ + -16.8517kJ   =  -261.7483kJ  + Hf

 

=>∆Hf  = -33.6kJmole-1.

 

Study the energy cycle diagram below and use it to:

 

(a)Identify the energy changes ∆H1 ∆H2 ∆H3 ∆H4 ∆H5 ∆H6

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

39

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

H1 – enthalpy/heat of formation of sodium chloride (∆Hf)
H2 -enthalpy/heat of atomization of sodium (∆Hat)
H3 -enthalpy/heat of ionization/ionization energy of sodium (∆H i)
H4enthalpy/heat of atomization of chlorine (∆Hat)
H5enthalpy/heat of electron affinity of chlorine (∆He)
H6 enthalpy/heat of lattice/Lattice energy of sodium chloride(∆H l)

 

  • Calculate ∆H1 given that ∆H2 =+108kJ , ∆H3=+500kJ, ∆H4 =+121kJ ,∆H5 =-364kJ and ∆H6 =-766kJ

 

Working:

H1 =∆H2 +∆H3 +∆H4 +∆H5 +∆H6

Substituting:

 

H1= +108kJ + +500kJ + +121kJ +-364kJ + -766kJH1= -401kJmole-1

 

  • Given the that:

 

  • Ionization energy of sodium = + 500kJmole-1 (ii)∆Hat of sodium = + 110kJmole-1

 

  • Electron affinity of chlorine = – 363kJmole-1 (iv)∆Hat of chlorine = + 120kJmole-1

 

(v) ∆Hf  of  sodium chloride = -411kJ , calculate the lattice energy of  sodium

chloride using an energy cycle diagram.

 

 

 

 

 

 

 

40

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Working:

Applying Hess law then:

Hf =∆Ha  +∆Hi +∆Ha +∆He +∆Hl

Substituting:

 

-411= +108kJ + +500kJ + +121kJ +-364kJ + x -411 + -108kJ + -500kJ + -121kJ + +364kJ = x

x= -776kJmole-1

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

41

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

42

 

20.0.0 REACTION RATES AND REVERSIBLE REACTIONS (15 LESSONS)

 

 

 

 

A.THE RATE OF CHEMICAL REACTION (CHEMICAL KINETICS)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

1.Introduction

 

The rate of a chemical reaction is the time taken for a given mass/amount of products to be formed. The rate of a chemical reaction is also the time taken for a given mass/amount of reactant to be consumed /used up.

 

Some reactions are too slow to be determined. e.g rusting ,decomposition of hydrogen peroxide and weathering.

 

Some reactions are too fast and instantaneous e.g. neutralization of acid and bases/alkalis in aqueous solution and double decomposition/precipitation.

 

Other reactions are explosive and very risky to carry out safely e.g. reaction of potassium with water and sodium with dilute acids.

 

The study of the rate of chemical reaction is useful in knowing the factors that influence the reaction so that efficiency and profitability is maximized in industries.

 

Theories of rates of reaction.

 

The rate of a chemical reaction is defined as the rate of change of concentration/amount of reactants in unit time. It is also the rate of formation of given concentration of products

 

in unit time.  i.e.

 

Rate of reaction  = Change in concentration/amount of reactants

 

Time taken for the change to occur

 

Rate of reaction = Change in concentration/amount of products formed Time taken for the products to form

 

For the above, therefore the rate of a chemical reaction is rate of decreasing reactants to form an increasing product.

The SI unit of time is second(s) but minutes and hours are also used.

 

(a)The collision theory

 

The collision theory is an application of the Kinetic Theory of matter which assumes matter is made up of small/tiny/minute particles like ions atoms and molecules. The collision theory proposes that

 

(i)for a reaction to occur, reacting particles must collide.

 

(ii)not all collisions between reacting particles are successful in a reaction. Collisions that initiate a chemical reaction are called successful / fruitful/ effective collisions

 

(iii)the speed at which particles collide is called collision frequency.

 

The higher the collision frequency the higher the chances of successful / fruitful/ effective collisions to form products.

 

(iv)the higher the chances of successful collisions, the faster the reaction.

 

(v)the average distance between solid particles from one another is too big for them to meet and collide successfully.

 

(vi)dissolving substances in a solvent ,make the solvent a medium for the reaction to take place.

 

The solute particle distance is reduced as the particle ions are free to move in the solvent medium.

 

(vii)successful collisions take place if the particles colliding have the required energy and right orientation which increases their vibration and intensity of successful / fruitful/ effective collisions to form products.

 

(b)The Activation Energy(Ea) theory

 

The Enthalpy of activation(∆Ha) /Activation Energy(Ea) is the minimum amount of energy which the reactants must overcome before they react. Activation Energy(Ea) is usually required /needed in bond breaking of the reacting particles. Bond breaking is an endothermic process that require an energy input.

 

The higher the bond energy the slower the reaction to start of.

Activation energy does not influence whether a reaction is exothermic or endothermic.

 

The energy level diagrams below shows the activation energy for exothermic and endothermic processes/reactions.

 

Energy level diagram showing the activation energy for exothermic processes /reactions.

 

Activated complex

 

  • B

 

A                                   B

Energy

 

 

kJ

 

Ea

 

A-A   B-B

 

 

 

A-B     A-B

 

 

 

 

Energy level diagram showing the activation energy for endothermic processes /reactions.

Reaction path/coordinate/path

 

 

Activated complex

 

  • B

 

A B
Energy
Ea
kJ A-B   A-B

 

 

∆Hr

A-A   B-B

 

 

 

 

Reaction path/coordinate/path

 

 

 

The activated complex is a mixture of many intermediate possible products which may not exist under normal physical conditions ,but can theoretically exist.

 

Exothermic reaction proceeds without further heating /external energy because it generates its own energy/heat to overcome activation energy.

 

Endothermic reaction cannot proceed without further heating /external energy because it does not generates its own energy/heat to overcome activation energy. It generally therefore requires continuous supply of more energy/heat to sustain it to completion.

 

  1. Measuring the rate of a chemical reaction.

 

The rate of a chemical reaction can be measure as:

(i)Volume of a gas in unit time;

  • if reaction is producing a gas as one of the products.

 

  • if reaction is using a gas as one reactants

 

(ii)Change in mass of reactants/products for solid products/reactants in unit time.

 

(iii)formation of a given mass of precipitate in unit time

 

(iv)a certain mass of reactants to completely form products/diminish.

 

Reactants may be homogenous or heterogenous.

-Homogenous reactions involve reactants in the same phase/state e.g.

solid-solid,gas-gas,liquid-liquid.

 

-Heterogenous reactions involve reactants in the different phase/state e.g. solid-liquid,gas-liquid,solid-gas.

 

  1. Factors influencing/altering/affecting/determining rate of reaction

 

The following factors alter/influence/affect/determine the rate of a chemical reaction:

 

(a)Concentration

(b)Pressure

 

  • Temperature (d)Surface area

 

(e)Catalyst

 

  1. Influence of concentration on rate of reaction

 

 

The higher the concentration, the higher the rate of a chemical reaction. An increase in concentration of the reactants reduces the distance between the reacting particles increasing their collision frequency to form products. Practically an increase in concentration reduces the time taken for the reaction to take place.

 

Practical determination of effect of concentration on reaction rate Method 1(a)

Reaction of sodium thisulphate with dilute hydrochloric acid

 

Procedure:

 

Measure 20cm3 of 0.05M sodium thisulphate into a 50cm3 glass beaker. Place the beaker on a white piece of filter paper with ink mark X on it. Measure 20cm3 of 0.1M hydrochloric acid solution using a 50cm3 measuring cylinder. Put the acid into the beaker containing sodium thisulphate. Immediately start off the stop watch/clock. Determine the time taken for the ink mark X to become invisible /obscured when viewed from above. Repeat the procedure by measuring different volumes of the acid and adding the volumes of the distilled water to complete table 1.

 

Sample results:Table 1.

 

Volume of Volume of Volume of Time taken for mark ‘X’ Reciprocal
acid(cm3) water(cm3) sodium to be of time
thiosulphate(cm3) invisible/obscured(seconds) 1
t
20.0 0.0 20.0 20.0 5.0 x 10-2
18.0 2.0 20.0 23.0 4.35 x 10-2
16.0 4.0 20.0 27.0 3.7 x 10-2
14.0 6.0 20.0 32.0 3.13 x 10-2
12.0 8.0 20.0 42.0 2.38 x 10-2
10.0 10.0 20.0 56.0 1.78 x 10-2

 

For most examining bodies/councils/boards the above results score for:

(a) complete table as evidence for all the practical work done and completed.

 

  • (i)Consistent use of a decimal point on time as evidence of understanding/knowledge of the degree of accuracy of stop watches/clock.

 

 

(ii)Consistent use of a minimum of four decimal points on inverse/reciprocal of time as evidence of understanding/knowledge of the degree of accuracy of scientific calculator.

 

  • accuracy against a school value based on candidate’s teachers-results

 

 

  • correct trend (time increase as more water is added/acid is diluted) in conformity with expected theoretical

 

Sample questions

 

  1. On separate graph papers plot a graph of:

(i)volume of acid used(x-axis) against time. Label this graph I

  • volume of acid used(x-axis) against 1/t. Label this graph II

 

  1. Explain the shape of graph I

Diluting/adding water is causes a decrease in concentration.

 

Decrease in concentration reduces the rate of reaction by increasing the time taken for reacting particle to collide to form products.

 

Sketch sample  Graph I

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Time

 

 

(seconds)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Volume of acid(cm3)

 

 

Sketch sample  Graph II

 

 

 

 

 

1/t

 

Sec-1  x 10-2

 

 

 

 

 

3.From graph II ,determine the time taken for the cross to be obscured/invisible when the volume of the acid is:

 

(i) 13cm3

From a correctly plotted graph

 

1/t  at 13cm3  on the graph => 2.75 x 10 -2
Volume of acid(cm3)
t  =  1 /  2.75 x 10-2   = 36.3636 seconds
(ii) 15cm3
From a correctly plotted graph
-2
Volume of acid (cm3)
1/t  at 15cm3  on the graph => 3.35 x 10
t  =  1 /  3.35 x 10-2   = 29.8507 seconds

 

(iii) 15cm3

 

From a correctly plotted graph

 

1/t at 17cm3 on the graph => 4.0 x 10-2 t = 1 / 4.0 x 10-2 = 25.0 seconds

 

(iv) 19cm3

 

From a correctly plotted graph

 

1/t  at 19cm3  on the graph => 4.65 x 10-2 t  =  1 /  4.65 x 10-2   =  21.5054 seconds

 

 

4.From graph II ,determine the volume of the acid used if the time taken for the cross to be obscured/invisible is:

 

(i)25 seconds

1/t  =>    1/25  = 4.0 x 10-2

Reading from a correctly plotted graph;

-2

 

(ii)30 seconds

1/t  =>    1/30  = 3.33 x 10-2

Reading from a correctly plotted graph;

-2

 

(iii)40 seconds

 

1/t  =>    1/40  = 2.5 x 10-2

Reading from a correctly plotted graph;

-2

 

 

  1. Write the equation for the reaction taking place

Na2S2O3 (aq) + 2HCl(aq) -> 2NaCl (aq)+ SO2 (g) + S(s) + H2O(l)

 

Ionically:

S2O32- (aq) + 2H+ (aq) ->  SO2 (g) + S(s) + H2O(l)

 

5.Name the yellow precipitate

 

Colloidal sulphur

 

 

 

Method 1(b)

 

Reaction of sodium thisulphate with dilute hydrochloric acid

 

You are provided with

2.0M Hydrochloric acid

0.4M sodium thiosulphate solution

 

Procedure:

 

Measure 10cm3 of sodium thisulphate into a 50cm3 glass beaker. Place the beaker on a white piece of filter paper with ink mark X on it.

 

Add 5.0cm3 of hydrochloric acid solution using a 10cm3 measuring cylinder into the beaker containing sodium thisulphate.

 

Immediately start off the stop watch/clock. Determine the time taken for the ink mark ‘Xto become invisible /obscured when viewed from above.

 

Repeat the procedure by measuring different volumes of the thiosulphate and adding the volumes of the distilled water to complete table 1.

 

Sample results:Table 1.

 

 

 

Volume Volume Volume of Concentation Time(T) taken T-1
of of water sodium of sodium for mark ‘X’ to
acid(cm3) (cm3) thiosulphate thisulphate in be invisible/
(cm3) molesdm-3 obscured(second
s)
5.0 0.0 25.0 0.4 20.0 5.0 x 10-2
5.0 5.0 20.0 0.32 23.0 4.35 x 10-2
5.0 10.0 15.0 0.24 27.0 3.7 x 10-2
5.0 15.0 10.0 0.16 32.0 3.13 x 10-2
Note concentration of diluted solution is got:
C1V1=C2V2 => 0.4 x 25 = C2x 25  =0.4M
C1V1=C2V2 => 0.4 x 20 = C2x 25  =0.32M
C1V1=C2V2 => 0.4 x 15 = C2x 25  =0.24M
C1V1=C2V2 => 0.4 x 10 = C2x 25  =0.16M

Sample questions

 

  1. On separate graph papers plot a graph of:

(i)Concentration of sodium thiosulphate against time. Label this graph I

 

(ii)Concentration of sodium thiosulphate against against T-1.Label this graph II

 

  1. Explain the shape of graph I

Diluting/adding water causes a decrease in concentration.

 

Decrease in concentration reduces the rate of reaction by increasing the time taken for reacting particle to collide to form products. From graph II

 

Determine the time taken if

 

(i)12cm3 of sodium thisulphate is diluted with 13cm3 of water. At 12cm3 concentration of sodium thisulphate

= C1V1=C2V2 => 0.4 x 1 2 = C2x 25  =0.192M
From correct graph at concentration 0.192M => 2.4 x10-2
I/t =  2.4 x10-2 t = 41.6667seconds
(ii)22cm3 of sodium thisulphate is diluted with 3cm3 of water.
At  22cm3 concentration of sodium thisulphate
= C1V1=C2V2 => 0.4 x 22 = C2x 25  =0.352M
From correct graph at concentration 0.352M => 3.6 x10-2
I/t =  3.6 x10-2 t = 27.7778seconds

 

Determine the volume of water and sodium thiosulphate if T-1 is 3.0 x10-1 From correct graph at T-1 = 3.0 x10-1 => concentration = 0.65 M

 

= C1V1=C2V2 => 0.4 x 25 = 0.65 M x V2 = 15.3846cm3 Volume of water = 25 – 15.3846cm3 = 9.6154cm3

 

Determine the concentration of hydrochloric acid if 12cm3 of sodium thiosulphate and 13cm3 of water was used.

 

At  12cm3 concentration of sodium thisulphate
= C1V1=C2V2 => 0.4 x 1 2 = C2x 25  =0.192M
Mole ratio Na2S2 O3 :HCl =1:2
Moles of Na2S2 O3 = 0.192M x 12 => 2.304 x 10 -3  moles
1000
Mole ratio  HCl =2.304 x 10-1  moles = 1.152 x 10-3  moles
2
Molarity o f HCl =  1.152 x 10-3  moles x 1000 = 0.2304M
5.0

 

Method 2

 

Reaction of Magnesium with dilute hydrochloric acid

Procedure

 

 

Scub 10centimeter length of magnesium ribbon with sand paper/steel wool. Measure 40cm3 of 0.5M dilute hydrochloric acid into a flask .Fill a graduated gas jar with water and invert it into a trough. Stopper the flask and set up the apparatus to collect the gas produced as in the set up below:

 

 

 

 

 

Hydrochloric acid

 

 

Hydrogen gas

 

Graduated gas jar

 

 

 

 

 

 

 

Magnesium ribbon

 

 

 

Carefully remove the stopper, carefully put the magnesium ribbon into the flask

 

  • cork tightly. Add the acid into the flask. Connect the delivery tube into the gas jar. Immediately start off the stop watch and determine the volume of the gas produced after every 30 seconds to complete table II below.

 

Sample results: Table II

 

Time(seconds) 0 30 60 90 120 150 180 210 240
Volume of gas 0.0 20.0 40.0 60.0 80.0 90.0 95.0 96.0 96.0
produced(cm3)

 

Sample practice questions

 

1.Plot a graph of volume of gas produced (y-axis) against time

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

2.Explain the shape of the graph.

The rate of reaction is faster when the concentration of the acid is high .

 

As time goes on, the concentration of the acid decreases and therefore less gas is produced.

 

When all the acid has reacted, no more gas is produced after 210 seconds and the graph flattens.

 

3.Calculate the rate of reaction at 120 seconds

 

From a tangent at 120 seconds rate of reaction = Change in volume of gas Change in time

=> From the tangent at 120seconds V 2 – V 1 = 96-84 = 12 = 0.2cm3sec-1
T2– T1 150-90 60

 

  1. Write an ionic equation for the reaction taking place. Mg2+(s) + 2H+(aq) -> Mg2+(aq) + H2 (g)

 

  1. On the same axis sketch then explain the curve that would be obtained if:
  • 1 M hydrochloric acid is used –Label this curve I (ii)1.0 M hydrochloric acid is used –Label this curve II

 

Observation:

Curve I is to the right

 

Curve II is to the left

 

Explanation

 

A decrease in concentration shift the rate of reaction graph to the right as more time is taken for completion of the reaction.

 

An increase in concentration shift the rate of reaction graph to the left as less time is taken for completion of the reaction.

Both graphs flatten after some time indicating the completion of the reaction.

 

b)Influence of pressure on rate of reaction

 

 

Pressure affects only gaseous reactants.

 

An increase in pressure reduces the volume(Boyles law) in which the particles are contained.

 

Decrease in volume of the container bring the reacting particles closer to each other which increases their chances of effective/successful/fruitful collision to form products.

 

An increase in pressure therefore increases the rate of reaction by reducing the time for reacting particles of gases to react.

 

At industrial level, the following are some reactions that are affected by pressure:

 

(a)Haber process for manufacture of ammonia N2(g) + 3H2(g) -> 2NH3(g)

 

(b)Contact process for manufacture of sulphuric(VI)acid 2SO2(g) + O2(g) -> 2SO3(g)

 

(c)Ostwalds process for the manufacture of nitric(V)acid

4NH3(g) + 5O2(g) -> 4NO (g) + 6H2O (l)

 

The influence of pressure on reaction rate is not felt in solids and liquids.

 

This is because the solid and liquid particles have fixed positions in their strong bonds and therefore no degree of freedom (Kinetic Theory of matter)

 

c)Influence of temperature on rate of reaction

 

An increase in temperature increases the kinetic energy of the reacting particles by increasing their collision frequency.

 

Increase in temperature increases the particles which can overcome the activation energy (Ea).

 

  • 10oC rise in temperature doubles the rate of reaction by reducing the time taken for the reaction to complete by a half.

 

Practical determination of effect of Temperature on reaction rate Method 1

 

Reaction of sodium thisulphate with dilute hydrochloric acid

 

Procedure:

 

Measure 20cm3 of 0.05M sodium thisulphate into a 50cm3 glass beaker. Place the beaker on a white piece of filter paper with ink mark X on it. Determine and record its temperature as room temperature in table 2 below. Measure 20cm3 of 0.1M hydrochloric acid solution using a 50cm3 measuring cylinder.

 

Put the acid into the beaker containing sodium thisulphate.

Immediately start off the stop watch/clock.

Determine the time taken for the ink mark X to become invisible /obscured when viewed from above.

 

Measure another 20cm3 separate portion of the thisulphate into a beaker, heat the solution to 30oC.

 

Add the acid into the beaker and repeat the procedure above. Complete table 2 below using different temperatures of the thiosulphate.

 

Sample results:Table 2.

 

Temperature of Na2S2O3 Room temperature 30 40 50 60
Time taken for mark X to 50.0 40.0 20.0 15.0 10.0
be obscured /invisible
(seconds)
Reciprocal of time(1/ ) 0.02 0.025 0.05 0.0667 0.1
t
Sample practice questions
  1. Plot a graph of temperature(x-axis) against 1/t

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

2(a)From your graph determine the temperature at which:

 

(i)1/t is ;

  1. 0.03

Reading directly from a correctly plotted graph = 32.25 oC

  1. 0.07

Reading directly from a correctly plotted graph = 48.0 oC

 

  • t is;

 

  1. 30 seconds

30 seconds => 1/t =1/30  =0.033

Reading directly from a correctly plotted graph 0.033 => 33.5 oC

 

  1. 45 seconds

45 seconds => 1/t =1/45  =0.022

Reading directly from a correctly plotted graph 0.022 => 29.0 oC

 

III. 25 seconds

25 seconds => 1/t =1/25  =0.04

Reading directly from a correctly plotted graph 0.04 => 36.0 oC

 

  • From your graph determine the time taken for the cross to become invisible at:

(i) 57.5 oC

 

Reading directly from a correctly plotted graph at 57.5 oC= 0.094 =>1/t = 0.094

t= 1/0.094  => 10.6383 seconds

 

(ii) 45 oC

Reading directly from a correctly plotted graph at 45 oC = 0.062 =>1/t = 0.062

 

t= 1/0.094  => 16.1290 seconds

 

(iii) 35 oC

Reading directly from a correctly plotted graph at 35 oC = 0.047 =>1/t = 0.047

 

t= 1/0.047  => 21.2766 seconds

 

Method 2

 

Reaction of Magnesium with dilute hydrochloric acid

Procedure

 

Scub 5centimeter length of magnesium ribbon with sand paper/steel wool.

Cut the piece into five equal one centimeter smaller pieces.

 

Measure 20cm3 of 1.0M dilute hydrochloric acid into a glass beaker .

Put one piece of the magnesium ribbon into the acid, swirl.

Immediately start off the stop watch/clock.

 

Determine the time taken for the effervescence/fizzing/bubbling to stop when viewed from above.

 

Record the time in table 2 at room temperature.

 

Measure another 20cm3 portions of 1.0M dilute hydrochloric acid into a clean beaker.

 

Heat separately one portion to 30oC, 40oC , 50oC and 60oC and adding 1cm length of the ribbon and determine the time taken for effervescence /fizzing /bubbling to stop when viewed from above .

 

Record each time to complete table 2 below using different temperatures of the acid.

 

Sample results:Table 1.

 

Temperature of acid(oC) Room temperature 30 40 50 60
Time taken effervescence 80.0 50.0 21.0 13.5 10.0
to stop (seconds)
Reciprocal of time(1/ ) 0.0125 0.02 0.0476 0.0741 0.1
t
Sample practice questions

 

  1. Plot a graph of temperature(x-axis) against 1/t

 

 

 

1/t

 

 

 

 

 

 

 

 

 

 

 

 

Temperature(oC)

 

2.(a)Calculate the number of moles of magnesium used given that 1cm of magnesium has a mass of 1g.(Mg= 24.0)

 

Moles = Mass of magnesium  => 1.0 4.167  x 10 -2 moles
Molar mass of Mg 24

 

(b)Calculate the number of moles of hydrochloric acid used Moles of acid = molarity x volume of acid

 

1000

-2

=> 1.0 x 20         = 2.0 x 10      moles

 

 

(c)Calculate the mass  of magnesium that remain unreacted

 

Mole ratio Mg: HCl = 1:2

Moles Mg = ½ moles HCl

 

=> ½ x 2.0 x 10 -2 moles = 1.0 x 10 -2 moles Mass of reacted Mg = moles x molar mass => 1.0 x 10 -2 moles x 24 = 0.24 g

 

Mass of unreacted Mg = Original total mass – Mass of reacted Mg => 1.0 g 0.24 = 0.76 g

 

(b)Calculate the total volume of hydrogen gas produced during the above reactions.

 

Mole ratio Mg : H2 = 1:1

Moles of Mg that reacted per experiment = moles H2 =1.0 x 10 -2 moles Volume of Hydrogen at s.t.p produced per experiment = moles x 24 dm3 => 1.0 x 10 -2 moles x 24 dm3 = 0.24dm3

Volume of Hydrogen at s.t.p produced  in 5 experiments =0.24 dm3 x 5

  • 2 dm3

 

3.(a)At what temperature was the time taken for magnesium to react equal to:

 

(i)70seconds

70 seconds => 1/t =1/70  =0.01429

Reading directly from a correctly plotted graph 0.01429 => 28.0 oC

 

(ii)40seconds

40 seconds => 1/t =1/40  =0.025

Reading directly from a correctly plotted graph 0.025 => 32.0 oC

 

(b)What is the time taken for magnesium to react if the reaction was done at:

(i) 55.0 oC

 

Reading directly from a correctly plotted graph at 55.0 oC=> 1/t = 8.0 x 10-2 => t = 1/8.0 x 10-2 = 12.5 seconds

 

(ii) 47.0 oC

Reading directly from a correctly plotted graph at 47.0 oC=> 1/t = 6.0 x 10-2 => t = 1/6.0 x 10-2 = 16.6667 seconds

 

(iii) 33.0 oC

Reading directly from a correctly plotted graph at 33.0 oC=> 1/t = 2.7 x 10-2 => t = 1/2.7 x 10-2 = 37.037 seconds

 

  1. Explain the shape of the graph.

 

Increase in temperature increases the rate of reaction as particles gain kinetic energy increasing their frequency and intensity of collision to form products.

 

d)Influence of surface area on rate of reaction

 

Surface area is the area of contact. An increase in surface area is a decrease in particle size. Practically an increase in surface area involves chopping /cutting

 

solid lumps into smaller pieces/chips then crushing the chips into powder. Chips thus have a higher surface area than solid lumps but powder has a highest surface area.

 

An increase in surface area of solids increases the area of contact with a liquid solution increasing the chances of successful/effective/fruitful collision to form products. The influence of surface area on rate of reaction is mainly in heterogeneous reactions.

 

Reaction of chalk/calcium carbonate on dilute hydrochloric acid Procedure

 

Measure 20cm3 of 1.0 M hydrochloric acid into three separate conical flasks labeled C1 C2 and C3 .

 

Using a watch glass weigh three separate 2.5g a piece of white chalk. Place the conical flask C1 on an electronic balance.

Reset the balance scale to 0.0.

Put one weighed sample of the chalk into the acid in the conical flask.

Determine the scale reading and record it at time =0.0.

 

Simultaneously start of the stop watch.

 

Determine and record the scale reading after every 30 seconds to complete Table I .

 

Repeat all the above procedure separately with C2 and C3 to complete Table II and Table III by cutting the chalk into small pieces/chips for C2 and crushing the chalk to powder for C3

Sample results:Table 1.

 

Time(seconds) 0.0 30.0 60.0 90.0 120.0 150.0 180.0 210.0 240.0
Mass of 2.5 2.0 1.8 1.4 1.2 1.0 0.8 0.5 0.5
CaCO3
Loss in mass 0.0 0.5 0.7 1.1 1.3 1.5 1.7 2.0 2.0
Sample results:Table 1I.
Time(seconds) 0.0 30.0 60.0 90.0 120.0 150.0 180.0 210.0 240.0
Mass of 2.5 1.9 1.5 1.3 1.0 0.8 0.5 0.5 0.5
CaCO3
Loss in mass 0.0 0.6 1.0 1.2 1.5 1.7 2.0 2.0 2.0

 

Sample results:Table III.

 

Time(seconds) 0.0 30.0 60.0 90.0 120.0 150.0 180.0 210.0 240.0
Mass of 2.5 1.8 1.4 1.0 0.8 0.5 0.5 0.5 0.5
CaCO3
Loss in mass 0.0 0.7 1.1 1.5 1.7 2.0 2.0 2.0 2.0

Sample questions:

 

1.Calculate the loss in mass made at the end of each time from the original to complete table I,II and III

 

2.On the same axes plot a graph of total loss in mass against time (x-axes) and label them curve I, II, and III from Table I, II, and III.

 

3.Explain why there is a loss in mass in all experiments.

 

Calcium carbonate react with the acid to form carbon(IV)oxide gas that escape to the atmosphere.

 

4.Write an ionic equation for the reaction that take place

CaCO3(s) + 2H+(aq) -> Ca2+(aq) + H2O(l) + CO2(g)

 

5.Sulphuric(VI)acid cannot be used in the above reaction. On the same axes sketch the curve which would be obtained if the reaction was attempted by reacting a piece of a lump of chalk with 0.5M sulphuric(VI)acid. Label it curve IV. Explain the shape of curve IV.

 

Calcium carbonate would react with dilute 0.5M sulphuric(VI)acid to form insoluble calcium sulphate(VI) that coat /cover unreacted Calcium carbonate stopping the reaction from reaching completion.

 

6.Calculate the volume of carbon(IV)oxide evolved(molar gas volume at room temperature = 24 dm3, C= 12.0, O= 16.O Ca=40.0)

 

Method I
Mole ratio  CaCO3(s) : CO2(g) = 1:1
Moles CaCO3(s) used = Mass CaCO 3 (s) 0.025 moles
Molar mass CaCO3(s)
Moles CO2(g) = 0.025 moles
Volume of  CO2(g) = moles x molar gas volume
=>0.025 moles  x  24 dm3 = 0.600 dm3/600cm3

 

Method II

Molar mass of CaCO3(s) = 100g produce 24 dm3 of CO2(g)

Mass of CaCO3(s) =2.5 g produce 2.5 x 24 = 0.600dm3 100

 

 

7.From curve I ,determine the rate of reaction (loss in mass per second)at time 180 seconds on the curve.

 

From  tangent at 180 seconds on curve I
Rate    = M 2 -M 1 => 2.08 1.375 = 0.625 = 0.006944g sec-1
T2– T1 222-132 90

 

8.What is the effect of particle size on the rate of reaction?

 

A larger surface area is a reduction in particle size which increases the area of contact between reacting particles increasing their collision frequency.

 

Theoretical examples

 

  1. Excess marble chips were put in a beaker containing 100cm3 of 0.2M hydrochloric acid. The beaker was then placed on a balance and total loss in mass recorded after every two minutes as in the table below.

 

Time(minutes) 0.0 2.0 4.0 6.0 8.0 10.0 12.0
Loss in mass(g) 0.0 1.80 2.45 2.95 3.20 3.25 3.25

 

(a)Why was there a loss in mass?

Carbon (IV) oxide gas was produced that escape to the surrounding

 

(b)Calculate the average rate of loss in mass between:

 

(i) 0 to 2 minutes 1.80 0.0 = 9.00g min-1
Average rate = M 2 -M 1 => = 1.8
2.0 0.0
T2– T1 2
(i) 6 to 8 minutes 3.20 2.95 = 0.125g min-1
Average rate = M 2 -M 1 => = 0.25
8.0 6.0
T2– T1 2

 

  • Explain the difference between the average rates of reaction in (i) and(ii) above.

 

Between 0 and 2 minutes , the concentration of marble chips and hydrochloric acid is high therefore there is a higher collision frequency between the reacting particles leading to high successful rate of formation of products.

 

Between 6 and 8 minutes , the concentration of marble chips and hydrochloric acid is low therefore there is low collision frequency between the reacting particles leading to less successful rate of formation of products.

 

(c)Write the equation for the reaction that takes place.

CaCO3(s) + 2HCl (aq) -> CaCO3 (aq) + H2O(l) + CO2(g)

 

(d)State and explain three ways in which the rate of reaction could be increased.

 

(i)Heating the acid- increasing the temperature of the reacting particles increases their kinetic energy and thus collision frequency.

 

(ii)Increasing the concentration of the acid-increasing in concentration reduces the distances between the reacting particles increasing their chances of effective/fruitful/successful collision to form products faster.

 

(iii)Crushing the marble chips to powder-this reduces the particle size/increase surface area increasing the area of contact between reacting particles.

 

(e)If the solution in the beaker was evaporated to dryness then left overnight in the open, explain what would happen.

 

It becomes wet because calcium (II) chloride absorbs water from the atmosphere and form solution/is deliquescent.

 

(f)When sodium sulphate (VI) was added to a portion of the contents in the beaker after the reaction , a white precipitate was formed .

(i)Name the white precipitate.

 

Calcium(II)sulphate(VI)

 

(ii)Write an ionic equation for the formation of the white precipitate

Ca2+(aq)  + SO42-(aq)->CaSO4(s)

 

(iii)State one use of the white precipitate

-Making plaster for building

-Manufacture of plaster of Paris

 

-Making sulphuric(VI)acid

 

(g)(i) Plot a graph of total loss in mass(y-axes) against time

 

(ii)From the graph, determine the rate of reaction at time 2 minutes. From a tangent/slope at 2 minutes;

Rate  of reaction = Average rate =M2-M1 => 2.25 1.30  = 0.95 = 0.3958g min-1
T2– T1 3.20 0.8 2.4

 

(iii)Sketch on the same axes the graph that would be obtained if 0.02M hydrochloric acid was used. Label it curve II

 

  1. Influence of catalyst on rate of reaction

 

 

Catalyst is a substance that alter the rate /speed of a chemical reaction but remain chemically unchanged at the end of a reaction. Biological catalysts are called enzymes. A catalyst does not alter the amount of products formed but itself may be altered physically e.g. from solid to powder to fine powder. Like biological enzymes, a catalyst only catalyse specific type of reactions

 

Most industrial catalysts are transition metals or their compounds. Catalyst works by lowering the Enthalpy of activation(∆H a)/activation energy (Ea) of the reactants .The catalyst lowers the Enthalpy of activation(∆Ha)/activation

energy (Ea) by:

 

  • forming short lived intermediate compounds called activated complex that break up to form the final product/s

 

  • being absorbed by the reactants thus providing the surface area on which reaction occurs.

 

A catalyst has no effect on the enthalpy of reaction Hr but only lowers the Enthalpy of activation(∆Ha)/activation energy (Ea)It thus do not affect/influence

 

whether the reaction is exothermic or endothermic as shown in the energy level diagrams below.

 

Energy level diagram showing the activation energy for exothermic processes /reactions.

 

Activated complex

 

  • B

 

A                                   B

Energy

 

 

kJ

 

 

Ea uncatalysed

 

A-A   B-B

Ea Catalysed

 

 

A-B     A-B

 

 

 

 

 

 

 

 

 

 

 

 

Energy level diagram showing the activation energy for endothermic processes /reactions.

 

Activated complex

 

  • B

 

A B
Energy
Ea
kJ A-B   A-B

 

 

∆Hr

A-A   B-B

 

 

 

 

Reaction path/coordinate/path

 

 

The following are some catalysed reaction processes.

 

(a)The contact process

Vanadium(V) Oxide(V2O5) or platinum(Pt) catalyses the oxidation of sulphur(IV)oxide during the manufacture of sulphuric(VI) acid from contact process.

SO2(g)          +             O2(g)  —-V2O5–>  SO3(g)

 

To reduce industrial cost of manufacture of sulphuric (VI) acid from contact process Vanadium(V) Oxide(V2O5) is used because it is cheaper though it is easily poisoned by impurities.

 

(b)Ostwalds process

 

Platinum promoted with Rhodium catalyses the oxidation of ammonia to nitrogen(II)oxide and water during the manufacture of nitric(V)acid

4NH3(g)   +           5O2(g)  —-Pt/Rh–>   4NO (g) + 6H2O(l)

 

(c)Haber process

 

Platinum or iron catalyses the combination of nitrogen and hydrogen to form ammonia gas

N2(g)  + 3H2(g)  —Pt or Fe—> 2NH3(g)

 

(d)Hydrogenation/Hardening of oil to fat

 

Nickel (Ni) catalyses the hydrogenation of unsaturated compound containing

  • C=C- or C=C- to saturated compounds without double or triple bond This process is used is used in hardening oil to fat.

 

(e)Decomposition of hydrogen peroxide Manganese(IV)oxide speeds up the rate of decomposition of hydrogen peroxide to water and oxygen gas.

 

This process/reaction is used in the school laboratory preparation of Oxygen. 2H2O2 (g) —-MnO2–> O2(g) + 2H2O(l)

 

(f)Reaction of metals with dilute sulphuric(VI)acid Copper(II)sulphate(VI) speeds up the rate of production of hydrogen gas from the reaction of Zinc and dilute sulphuric(VI)acid.

 

This process/reaction is used in the school laboratory preparation of Hydrogen.

 

H2 SO4 (aq) +  Zn(s)  —-CuSO4–>      ZnSO4 (aq)  +  H2(g)

 

(g) Substitution reactions

 

When placed in bright sunlight or U.V /ultraviolet light , a mixture of a halogen and an alkane undergo substitution reactions explosively to form halogenoalkanes. When paced in diffused sunlight the reaction is very slow.

 

e.g.  CH4(g)  +    Cl2(g)  —u.v. light–>  CH3Cl(g)  +   HCl(g)

 

(h)Photosynthesis

 

Plants convert carbon(IV)oxide gas from the atmosphere and water from the soil to form glucose and oxygen as a byproduct using sunlight / ultraviolet light.

 

6CO2(g)   +    6H2O(l)  —u.v. light–> C6H12O6(g) +     O2(g)

 

(i)Photography

 

Photographic film contains silver bromide emulsion which decomposes to silver and bromine on exposure to sunlight.

2AgBr(s)  —u.v/sun light–> 2Ag(s)  + Br2(l)

 

When developed, the silver deposits give the picture of the object whose photograph was taken depending on intensity of light. A picture photographed in diffused light is therefore blurred.

 

 

 

Practical determination of effect of catalyst on decomposition of hydrogen peroxide

 

Measure 5cm3 of 20 volume hydrogen peroxide and then dilute to make 40cm3 in a measuring cylinder by adding distilled water.

 

Divide it into two equal portions.

 

(i)Transfer one 20cm3volume hydrogen peroxide into a conical/round bottomed/flat bottomed flask. Cork and swirl for 2 minutes. Remove the cork. Test the gas produced using a glowing splint. Clean the conical/round bottomed/flat bottomed flask.

 

(ii)Put 2.0g of Manganese (IV) oxide into the clean conical/round bottomed/flat bottomed flask. Stopper the flask.

 

Transfer the second portion of the 20cm3volume hydrogen peroxide into a conical/round bottomed/flat bottomed flask through the dropping/thistle funnel. Connect the delivery tube to a calibrated/graduated gas jar as in the set up below.

 

Start off the stop watch and determine the volume of gas in the calibrated/graduated gas jar after every 30 seconds to complete Table 1.

 

(iii)Weigh a filter paper .Use the filter paper to filter the contents of the conical conical/round bottomed/flat bottomed flask. Put the residue on a sand bath to dry. Weigh the dry filter paper again .Determine the new mass Manganese (IV) oxide.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Time(seconds) 0.0 30.0 60.0 90.0 120.0 150.0 180.0 210.0 240.0 270.0
Volume of gas 0.0 20.0 40.0 60.0 80.0 90.0 95.0 96.0 96.0 96.0
(cm3)
Mass of MnO2 before reaction(g) Mass of MnO2 after reaction(g)
2.0 2.0

 

Plot a graph of volume of gas produced against time(x-axes)

 

Catalysed reaction

 

 

Uncatalysed reaction

 

 

 

 

  1. b) On the same axes, plot a graph of the uncatalysed reaction.
  • Explain the changes in mass of manganese(IV)oxide before and after the

reaction.

The mass of MnO2 before and after the reaction is the same but a more fine powder after the experiment. A catalyst therefore remains unchanged chemically but may physically change.

 

 

B.EQUILIBRIA (CHEMICAL CYBERNETICS)

 

Equilibrium is a state of balance.

 

Chemical equilibrium is state of balance between the reactants and products. As reactants form products, some products form back the reactants. Reactions in which the reactants form products to completion are said to be reversible i.e.

 

A          +             B         ->           C         +             D

 

Reactions in which the reactants form products and the products can reform the reactants are said to be reversible.

 

A+B C+D Reversible reactions may be:

 

(a)Reversible physical changes

 

(b)Reversible chemical changes

(c)Dynamic equilibrium

 

(a)Reversible physical changes

 

Reversible physical change is one which involves:

 

  • change of state/phase from solid, liquid, gas or aqueous solutions. States of matter are interconvertible and a reaction involving a change from one state/phase can be reversed back to the original.

 

  • colour changes. Some substances/compounds change their colours without change in chemical substance.

 

Examples of reversible physical changes

 

  • colour change on heating and cooling:

 

  1. Zinc(II)Oxide changes from white when cool/cold to yellow when

hot/heated and back.

ZnO(s)    ZnO(s)

(white when cold)                     (yellow when hot)

 

  1. Lead(II)Oxide changes from yellow when cold/cool to brown when

 

hot/heated and back.

PbO(s)      PbO(s)

 

(brown when hot)                      (yellow when cold)

 

(ii)Sublimation

 

  1. Iodine sublimes from a grey crystalline solid on heating to purple vapour. Purple vapour undergoes deposition back to the grey crystalline solid.

 

I2(s)   I2(g)

(grey crystalline solid              (purple vapour

undergo sublimation)               undergo deposition)

 

  1. Carbon (IV)oxide gas undergoes deposition from a colourless gas to a white solid at very high pressures in a cylinder. It sublimes back to the colourless gas if pressure is reduced

 

CO2(s)                                            CO2(g)

(white powdery solid               (colourless/odourless gas

 

undergo sublimation)               undergo deposition)

 

(iii)Melting/ freezing and boiling/condensation

 

Ice on heating undergo melting to form a liquid/water. Liquid/water on further heating boil/vaporizes to form gas/water vapour. Gas/water vapour on cooling, condenses/liquidifies to water/liquid. On further cooling, liquid water freezes to ice/solid.

 

Melting boiling
H2O(s) H2O(l) H2O(s)
Freezing condensing

 

(iv)Dissolving/ crystallization/distillation

 

Solid crystals of soluble substances (solutes) dissolve in water /solvents to form a uniform mixture of the solute and solvent/solution. On crystallization /distillation /evaporation the solvent evaporate leaving a solute back. e.g.

NaCl(s)        +  aq      NaCl(aq)

 

(b)Reversible chemical changes

 

These are reactions that involve a chemical change of the reactants which can be reversed back by recombining the new substance formed/products.

 

Examples of Reversible chemical changes

 

(i)Heating Hydrated salts/adding water to anhydrous salts.

 

When hydrated salts are heated they lose some/all their water of crystallization and become anhydrous.Heating an unknown substance /compound that forms a colourless liquid droplets on the cooler parts of a dry test/boiling tube is in fact

 

  • confirmation inference that the substance/compound being heated is hydrated.

 

When anhydrous salts are added (back) some water they form hydrated compound/salts.

 

Heating Copper(II)sulphate(VI)pentahydrate and cobalt(II)chloride hexahydrate (i)Heat about 5.0g of Copper(II)sulphate(VI) pentahydrate in a clean dry

 

test tube until there is no further colour change on a small Bunsen flame. Observe any changes on the side of the test/boiling tube. Allow the boiling tube to cool.Add about 10 drops of distilled water. Observe any changes.

 

(ii)Dip a filter paper in a solution of cobalt(II)chloride hexahydrate. Pass one end the filter paper to a small Bunsen flame repeatedly. Observe any changes on the filter paper. Dip the paper in a beaker containing distilled water. Observe any changes.

 

Sample observations

Hydrated Observation Observation after heating Observation on
compound before heating adding water
Copper(II)sulphate Blue crystalline (i)colour changes from blue (i)colour changes
(VI) pentahydrate solid to white. from white to blue
(ii)colourless liquid forms on (ii)boiling tube
the cooler parts of boiling / becomes warm /hot.
test tube
Cobalt(II)chloride Pink crystalline (i)colour changes from pink (i)colour changes
hexahydrate solid/solution to blue. from blue to pink
(ii) colourless liquid forms on (ii)boiling tube
the cooler parts of boiling / becomes warm/hot.
test tube (if crystal are used)

 

When blue Copper(II)sulphate (VI) pentahydrate is heated, it loses the five molecules of water of crystallization to form white anhydrous Copper(II)sulphate (VI).Water of crystallization form and condenses as colourless droplets on the cooler parts of a dry boiling/test tube.

 

This is a chemical change that produces a new substance. On adding drops of water to an anhydrous white copper(II)sulphate(VI) the hydrated compound is

 

formed back. The change from hydrated to anhydrous and back is therefore reversible chemical change.Both anhydrous white copper(II)sulphate(VI) and blue cobalt(II)chloride hexahydrate are therefore used to test for the presence of water when they turn to blue and pink respectively.

CuSO4(s)+ 5H2 O(l) CuSO4 .5H2 O(s/aq)
(white/anhydrous) (blue/hydrated)
CoCl2(s)+ 6H2 O(l) CoCl2 .6H2 O(s/aq)
(blue/anhydrous) (pink/hydrated)

 

(ii)Chemical sublimation

 

Some compounds sublime from solid to gas by dissociating into new different compounds. e.g.

 

Heating ammonium chloride

 

(i)Dip a glass rod containing concentrated hydrochloric acid. Bring it near the mouth of a bottle containing concentrated ammonia solution. Explain the observations made.

 

When a glass rod containing hydrogen chloride gas is placed near ammonia gas, they react to form ammonium chloride solid that appear as white fumes.

 

This experiment is used interchangeably to test for the presence of hydrogen chloride gas (and hence Cl ions) and ammonia gas (and hence NH4+ ions)

 

(ii)Put 2.0 g of ammonium chloride in a long dry boiling tube. Place wet / moist /damp blue and red litmus papers separately on the sides of the mouth of the boiling tube. Heat the boiling tube gently then strongly. Explain the observations made.

 

When ammonium chloride is heated it dissociates into ammonia and hydrogen chloride gases. Since ammonia is less dense, it diffuses faster to turn both litmus papers blue before hydrogen chloride turn red because it is denser. The heating and cooling of ammonium chloride is therefore a reversible chemical change.

NH3(g)                 +              HCl(g)        NH4Cl(s)

(Turns moist                        (Turns moist                  (forms white fumes)

 

litmus paper blue)           litmus paper red)

 

 

 

 

(c)Dynamic equilibria

 

For reversible reactions in a closed system:

 

(i) at the beginning;

 

-the reactants are decreasing in concentration with time -the products are increasing in concentration with time

 

  • after some time a point is reached when as the reactants are forming products the products are forming reactants. This is called equilibrium. Sketch showing the changes in concentration of reactants and products in a

 

 

closed system

Reactants concentration

 

decreases to form

 

products

 

 

 

 

Concentration

 

 

 

 

 

 

Equilibrium established /rate of formation of products equal to rate of formation of reactants.

 

 

Mole dm-3

Products concentration

 

increases from time=0.0

 

 

 

Reaction progress/path/coordinate

 

For a system in equilibrium:

 

  • a reaction from left to right (reactants to products) is called forward reaction.

 

  • a reaction from right to left (products to reactants) is called backward

reaction.

 

(iii)a reaction in which the rate of forward reaction is equal to the rate of backward reaction is called a dynamic equilibrium.

 

A dynamic equilibrium is therefore a balance of the rate of formation of products and reactants. This balance continues until the reactants or products are disturbed/changed/ altered.

 

The influence of different factors on a dynamic equilibrium was first investigated from 1850-1936 by the French Chemist Louis Henry Le Chatellier. His findings were called Le Chatelliers Principle which states that:

 

“if a stress/change is applied to a system in dynamic equilibrium, the system readjust/shift/move/behave so as to remove/ reduce/ counteract/ oppose the stress/change”

 

Le Chatelliers Principle is applied in determining the effect/influence of several factors on systems in dynamic equilibrium. The following are the main factors that influence /alter/ affect systems in dynamic equilibrium:

 

(a)Concentration

(b)Pressure

 

(c)Temperature

(d)Catalyst

 

(a)Influence of concentration on dynamic equilibrium

 

An increase/decrease in concentration of reactants/products at equilibrium is a stress. From Le Chatelliers principle the system redjust so as to remove/add the excessreduced concentration.

 

Examples of influence of concentration on dynamic equilibrium (i)Chromate(VI)/CrO42- ions in solution are yellow. Dichromate(VI)/Cr2O72-ions in solution are orange. The two solutions exist in equilibrium as in the equation:

 

2H+ (aq)  +  2CrO 2- (aq) Cr O 2- (aq) + H O(l)
4 7
2 2
(Yellow) (Orange)

 

  1. I. If an acid is/H+ (aq) is added to the equilibrium mixture a stress is created on the reactant side where there is already H+ The equilibrium shift forward to the right to remove/reduce the excess H+ ions added. Solution mixture becomes More Cr2O72- ions formed in the solution mixture make it to be more orange in colour.

 

  1. II. If a base/OH (aq) is added to the equilibrium mixture a stress is created on the reactant side on the H+ H+ ions react with OH (aq) to form water.

 

H+ (aq) +OH (aq) -> H2O(l)

 

The equilibrium shift backward to the left to add/replace the H+ ions that have reacted with the OH (aq) ions . More of the CrO42- ions formed in the solution mixture makes it to be more yellow in colour.

 

2OH (aq)  +  2Cr2O7 2- (aq) CrO4 2-  (aq)   +  H2O(l)
(Orange) (Yellow)

 

  1. I. If an acid/ H+ (aq) is added to the equilibrium mixture a stress is created on the reactant side on the OH (aq). H+ ions react with OH (aq) to form water.

H+ (aq) +OH (aq) -> H2O(l)

 

The equilibrium shift backward to the left to add/replace the 2OH (aq) that have reacted with the H+ (aq) ions . More Cr2O72- (aq)ions formed in the solution mixture makes it to be more Orange in colour.

 

  1. II. If a base /OH (aq) is added to the equilibrium mixture a stress is created on the reactant side where there is already OH (aq) ions. The equilibrium shift forward to the right to remove/reduce the excess OH (aq) ions added. More of the Cr2O72- ions are formed in the solution mixture making it to be more orange in colour.

 

(i)Practical determination of the influence of alkali/acid on Cr2O72- / CrO42- equilibrium mixture

 

 

Measure about 2 cm3 of Potassium dichromate (VI) solution into a test tube.

 

Note that the solution mixture is orange.

Add three drops of 2M sulphuric(VI) acid. Shake the mixture carefully.

 

Note that the solution mixture is remains orange.

Add about six drops of 2M sodium hydroxide solution. Shake carefully.

 

Note that the solution mixture is turns yellow.

Explanation

 

The above observations can be explained from the fact that both the dichromate(VI)and chromate(VI) exist in equilibrium. Dichromate(VI) ions are stable in acidic solutions while chromate(VI)ions are stable in basic solutions. An equilibrium exist thus:

 

OH-

Cr2O72-     CrO42-

H+

 

When an acid is added, the equilibrium shift forward to the right and the mixture become more orange as more Cr2O72- ions exist.

 

When a base is added, the equilibrium shift backward to the left and the mixture become more yellow as more CrO42- ions exist.

 

(ii)Practical determination of the influence of alkali/acid on bromine water in an equilibrium mixture

 

Measure 2cm3 of bromine water into a boiling tube. Note its colour.

 

Bromine water is yellow

Add three drops of 2M sulphuric(VI)acid. Note  any colour change

 

Colour becomes more yellow

 

Add seven drops of 2M sodium hydroxide solution. Note any colour change.

 

Solution mixture becomes colourless/Bromine water is decolourized. Explanation

 

When added distilled water,an equilibrium exist between bromine liquid (Br2(aq)) and the bromide ion(Br), hydrobromite ion(OBr) and hydrogen ion(H+) as in the equation:

 

H2O(l)  + Br2(aq)  OBr (aq) +  H+ (aq)  +  Br (aq)

If an acid (H+)ions is added to the equilibrium mixture, it increases the concentration of the ions on the product side which shift backwards to the left to remove the excess H+ ions on the product side making the colour of the solution mixture more yellow.

 

If a base/alkali OH is added to the equilibrium mixture, it reacts with H+ ions on the product side to form water.

H+ (aq)+ OH(aq) -> H2O(l)

 

This decreases the concentration of the H+ ions on the product side which shift the equilibrium forward to the right to replace H+ ions making the solution mixture colourless/less yellow (Bromine water is decolorized)

 

(iii)Practical determination of the influence of alkali/acid on common acid-base indicators.

 

Place 2cm3 of phenolphthalein ,methyl orange and litmus solutions each in three separate test tubes.

 

To each test tube add two drops of water. Record your observations in Table 1 below.

 

To the same test tubes, add three drops of 2M sulphuric(VI)acid. Record your observations in Table 1 below.

 

To the same test tubes, add seven drops of 2M sodium hydroxide solution.

Record your observations in Table 1 below.

To the same test tubes, repeat adding four drops of 2M sulphuric(VI)acid.

 

Table 1
Indicator Colour of indicator in
Water Acid(2M sulphuric Base(2M sodium
(VI) acid) hydroxide)
Phenolphthalein Colourless Colourless Pink
Methyl orange Yellow Red Orange
Litmus solution Colourless Red Blue

 

Explanation

 

An indicator is a substance which shows whether another substance is an acid , base or neutral.

 

Most indicators can be regarded as very weak acids that are partially dissociated into ions.An equilibrium exist between the undissociated molecules and the dissociated anions. Both the molecules and anions are coloured. i.e.

 

HIn(aq) H+ (aq)  +  In (aq)
(undissociated indicator (dissociated indicator
molecule(coloured)) molecule(coloured))

When an acid H+ is added to an indicator, the H+ ions increase and equilibrium shift backward to remove excess H+ ions and therefore the colour of the undissociated (HIn) molecule shows/appears.

 

When a base/alkali OH is added to the indicator, the OH reacts with H+ ions from the dissociated indicator to form water.

H+ (aq)         +           OH-(aq)        -> H2O(l)

(from indicator)     (from alkali/base)

 

The equilibrium shift forward to the right to replace the H+ ion and therefore the colour of dissociated (In) molecule shows/appears.

 

The following examples illustrate the above.
(i)Phenolphthalein indicator exist as:
HPh H+ (aq) + Ph(aq)
(colourless molecule) (Pink anion)

On adding an acid ,equilibrium shift backward to the left to remove excess H+ ions and the solution mixture is therefore colourless.

 

When a base/alkali OH is added to the indicator, the OH reacts with H+ ions from the dissociated indicator to form water.

H+ (aq)         +           OH-(aq)        -> H2O(l)

 

(from indicator)     (from alkali/base)

 

The equilibrium shift forward to the right to replace the removed/reduced H+ ions. The pink colour of dissociated (Ph) molecule shows/appears.

(ii)Methyl Orange indicator exists as:
HMe H+ (aq) + Me(aq)
(Red molecule) (Yellow/Orange anion)

On adding an acid ,equilibrium shift backward to the left to remove excess H+ ions and the solution mixture is therefore red.

 

When a base/alkali OH is added to the indicator, the OH reacts with H+ ions from the dissociated indicator to form water.

H+ (aq)         +           OH-(aq)        -> H2O(l)

(from indicator)     (from alkali/base)

 

The equilibrium shift forward to the right to replace the removed/reduced H+ ions. The Orange colour of dissociated (Me) molecule shows/appears.

 

(b)Influence of Pressure on dynamic equilibrium

 

Pressure affects gaseous reactants/products. Increase in pressure shift/favours the equilibrium towards the side with less volume/molecules. Decrease in pressure shift the equilibrium towards the side with more volume/molecules. More yield of products is obtained if high pressures produce less molecules / volume of products are formed.

 

If the products and reactants have equal volume/molecules then pressure has no effect on the position of equilibrium

 

The following examples show the influence of pressure on dynamic equilibrium:

 

(i)Nitrogen(IV)oxide /Dinitrogen tetroxide mixture

 

Nitrogen(IV)oxide and dinitrogen tetraoxide can exist in dynamic equilibrium in a closed test tube. Nitrogen(IV)oxide is a brown gas. Dinitrogen tetraoxide is a yellow gas.

Chemical equation : 2NO2(g) = ==== N2 O4 (g)
Gay Lussacs law 2Volume 1Volume
Avogadros law 2molecule 1molecule

 

 

2 volumes/molecules of Nitrogen(IV)oxide form 1 volumes/molecules of dinitrogen tetraoxide

 

Increase in pressure shift the equilibrium forward to the left where there is less volume/molecules.The equilibrium mixture become more yellow.

 

Decrease in pressure shift the equilibrium backward to the right where there is more volume/molecules. The equilibrium mixture become more brown.

 

(ii)Iodine vapour-Hydrogen gas/Hydrogen Iodide mixture.

Pure hydrogen gas reacts with Iodine vapour to form Hydrogen Iodide gas.

Chemical equation : I2(g) + H2(g) = ==== 2HI (g)
Gay Lussacs law 1Volume 1Volume 2Volume
Avogadros law 1molecule 1molecule 2molecule

 

 

(1+1) 2 volumes/molecules of Iodine and Hydrogen gasform 2 volumes/molecules of Hydrogen Iodide gas.

 

Change in pressure thus has no effect on position of equilibrium.

 

(iii)Haber process.

 

Increase in pressure of the Nitrogen/Hydrogen mixture favours the formation of more molecules of Ammonia gas in Haber process.

 

The yield of ammonia is thus favoured by high pressures

 

 

Chemical equation :

 

 

N2(g)

 

 

+

 

 

3H2

 

 

(g)

 

 

->

 

 

2NH3 (g)

 

 

 

Gay Lussacs law Avogadros law

 

1Volume

1molecule

 

3Volume

3molecule

 

2Volume

2molecule

 

 

 

(1 + 3) 4 volumes/molecules of Nitrogen and Hydrogen react to form 2 volumes/molecules of ammonia.

 

Increase in pressure shift the equilibrium forward to the left where there is less volume/molecules.

 

The yield of ammonia increase.

 

Decrease in pressure shift the equilibrium backward to the right where there is more volume/molecules.

 

The yield of ammonia decrease.

 

(iv)Contact  process.

 

Increase in pressure of the Sulphur(IV)oxide/Oxygen mixture favours the formation of more molecules of Sulphur(VI)oxide gas in Contact process. The yield of Sulphur(VI)oxide gas is thus favoured by high pressures.

 

Chemical equation : 2SO2(g)+ O2 (g)-> 2SO3 (g)
Gay Lussacs law 2Volume 1Volume 2Volume
Avogadros law 2molecule 1molecule 2molecule

 

 

(2 + 1) 3 volumes/molecules of Sulphur(IV)oxide/Oxygen mixture react to form 2 volumes/molecules of Sulphur(VI)oxide gas.

 

Increase in pressure shift the equilibrium forward to the left where there is less volume/molecules. The yield of Sulphur(VI)oxide gas increase.

 

Decrease in pressure shift the equilibrium backward to the right where there is more volume/molecules. The yield of Sulphur(VI)oxide gas decrease.

 

(v)Ostwalds process.

 

Increase in pressure of the Ammonia/Oxygen mixture favours the formation of more molecules of Nitrogen(II)oxide gas and water vapour in Ostwalds process. The yield of Nitrogen(II)oxide gas and water vapour is thus favoured by low pressures.

 

Chemical equation : 4NH3(g) +  5O2 (g) ->  4NO(g) +   6H2O (g)
Gay Lussacs law 4Volume 5Volume 4Volume 6Volume
Avogadros law 4molecule 5molecule 4molecule 6Molecule

 

 

(4 + 5) 9 volumes/molecules of Ammonia/Oxygen mixture react to form 10 volumes/molecules of Nitrogen(II)oxide gas and water vapour.

 

Increase in pressure shift the equilibrium backward to the left where there is less volume/molecules. The yield of Nitrogen(II)oxide gas and water vapour decrease.

 

Decrease in pressure shift the equilibrium forward to the right where there is more volume/molecules. The yield of Nitrogen(II)oxide gas and water vapour increase.

 

Note
If the water vapour is condensed on cooling, then:
6H2O (l)
Chemical equation : 4NH3(g) +  5O2 (g) ->  4NO(g)    +
Gay Lussacs law 4Volume 5Volume 4Volume 0Volume
Avogadros law 4molecule 5molecule 4molecule 0Molecule

 

(4 + 5) 9 volumes/molecules of Ammonia/Oxygen mixture react to form 4 volumes/molecules of Nitrogen(II)oxide gas and no vapour.

 

Increase in pressure shift the equilibrium forward to the right where there is less volume/molecules. The yield of Nitrogen(II)oxide gas increase.

 

Decrease in pressure shift the equilibrium backward to the left where there is more volume/molecules. The yield of Nitrogen(II)oxide gas decrease.

 

(c)Influence of Temperature on dynamic equilibrium

 

 

  • decrease in temperature favours the reaction that liberate/generate more heat thus exothermic reaction(-ΔH).

 

An increase in temperature favours the reaction that do not liberate /generate more heat thus endothermic reaction(+ΔH).

 

Endothermic reaction are thus favoured by high temperature/heating Exothermic reaction are favoured by low temperature/cooling.

 

If a reaction/equilibrium mixture is neither exothermic or endothermic, then a change in temperature/cooling/heating has no effect on the equilibrium position.

 

(i)Nitrogen(IV)oxide /Dinitrogen tetroxide mixture

 

 

Nitrogen(IV)oxide and dinitrogen tetraoxide can exist in dynamic equilibrium in a closed test tube. Nitrogen(IV)oxide is a brown gas. Dinitrogen tetraoxide is a yellow gas.

 

Chemical equation :                   2NO2(g)     =====                 N2 O4 (g)

 

On heating /increasing temperature, the mixture becomes more brown. On cooling the mixture become more yellow. This show that

 

(i)the forward reaction to the right is exothermic(-ΔH).

 

On heating an exothermic process the equilibrium shifts to the side that generate /liberate less heat.

 

(ii)the backward reaction to the right is endothermic(+ΔH).

 

On cooling an endothermic process the equilibrium shifts to the side that do not generate /liberate heat.

 

(c)Influence of Catalyst on dynamic equilibrium

 

 

A catalyst has no effect on the position of equilibrium. It only speeds up the rate of attainment. e.g.

 

Esterification of alkanols and alkanoic acids naturally take place in fruits.In the laboratory concentrated sulphuric(VI)acid catalyse the reaction.The equilibrium mixture forms the ester faster but the yield does not increase.

CH3CH2OH(l)+CH3COOH(l) ==Conc.H2SO4== CH3COOCH2CH3(aq) + H2O(l)

 

 

 

(d)Influence of rate of reaction and dynamic equilibrium (Optimum conditions) on industrial processes

 

 

 

 

Industrial processes are commercial profit oriented. All industrial processes take place in closed systems and thus in dynamic equilibrium.

 

For manufacturers, obtaining the highest yield at minimum cost and shortest time is paramount.

 

The conditions required to obtain the highest yield of products within the shortest time at minimum cost are called optimum conditions

Optimum condition thus require understanding the effect of various factors on:

 

(i)rate of reaction(Chemical kinetics)

 

(ii)dynamic equilibrium(Chemical cybernetics)

 

1.Optimum condition in Haber process Chemical equation

 

N2 (g)  +  3H2 (g)  ===Fe/Pt===  2NH3 (g)      ΔH = -92kJ

 

Equilibrium/Reaction rate considerations

 

(i)Removing ammonia gas once formed shift the equilibrium forward to the right to replace the ammonia. More/higher yield of ammonia is attained.

 

(ii)Increase in pressure shift the equilibrium forward to the right where there is less volume/molecules . More/higher yield of ammonia is attained. Very high pressures raises the cost of production because they are expensive to produce and maintain. An optimum pressure of about 500atmospheres is normally used.

 

 

(iii)Increase in temperature shift the equilibrium backward to the left because the reaction is exothermic(ΔH = -92kJ) . Ammonia formed decomposes back to Nitrogen and Hydrogen to remove excess heat therefore a less yield of ammonia is attained. Very low temperature decrease the collision frequency of Nitrogen and Hydrogen and thus the rate of reaction too slow and uneconomical.

An optimum temperature of about 450oC is normally used.

 

 

(iv)Iron and platinum can be used as catalyst. Platinum is a better catalyst but more expensive and easily poisoned by impurities than Iron. Iron is promoted /impregnated with AluminiumOxide(Al2O3) to increase its surface area/area of contact with reactants and thus efficiency.The catalyst does not increase the yield of ammonia but it speed up its rate of formation.

 

2.Optimum condition in Contact process

 

Chemical equation

 

2SO2 (g) + O2 (g) ===V2O5/Pt=== 2SO3 (g) ΔH = -197kJ Equilibrium/Reaction rate considerations

 

(i)Removing sulphur(VI)oxide gas once formed shift the equilibrium forward to the right to replace the sulphur(VI)oxide. More/higher yield of sulphur(VI) oxide is attained.

 

(ii)Increase in pressure shift the equilibrium forward to the right where there is less volume/molecules . More/higher yield of sulphur(VI)oxide is attained. Very high pressures raises the cost of production because they are expensive to produce and maintain. An optimum pressure of about 1-2 atmospheres is normally used to attain about 96% yield of SO3.

 

 

(iii)Increase in temperature shift the equilibrium backward to the left because the reaction is exothermic(ΔH = -197kJ) . Sulphur(VI)oxide formed decomposes back to Sulphur(IV)oxide and Oxygen to remove excess heat therefore a less yield of Sulphur(VI)oxide is attained. Very low temperature decrease the collision frequency of Sulphur(IV)oxide and Oxygen and thus the rate of reaction too slow and uneconomical.

An optimum temperature of about 450oC is normally used.

 

 

(iv)Vanadium(V)Oxide and platinum can be used as catalyst. Platinum is a better catalyst and less easily poisoned by impurities but more expensive. Vanadium(V)Oxide is very cheap even if it is easily poisoned by impurities. The catalyst does not increase the yield of Sulphur (VI)Oxide but it speed up its rate of formation.

 

 

3.Optimum condition in Ostwalds process

 

Chemical equation

 

4NH3 (g)  +  5O2 (g) ===Pt/Rh===  4NO (g) + 6H2O (g) ΔH = -950kJ

 

Equilibrium/Reaction rate considerations

 

(i)Removing Nitrogen(II)oxide gas once formed shift the equilibrium forward to the right to replace the Nitrogen(II)oxide. More/higher yield of Nitrogen(II) oxide is attained.

 

(ii)Increase in pressure shift the equilibrium backward to the left where there is less volume/molecules . Less/lower yield of Nitrogen(II)oxide is attained.

 

Very low pressures increases the distance between reacting NH3and O2 molecules.

 

An optimum pressure of about 9 atmospheres is normally used.

 

 

(iii)Increase in temperature shift the equilibrium backward to the left because the reaction is exothermic(ΔH = -950kJ) . Nitrogen(II)oxide and water vapour formed decomposes back to Ammonia and Oxygen to remove excess heat therefore a less yield of Nitrogen(II)oxide is attained. Very low temperature decrease the collision frequency of Ammonia and Oxygen and thus the rate of reaction too slow and uneconomical.

An optimum temperature of about 900oC is normally used.

 

(iv)Platinum can be used as catalyst. Platinum is very expensive.It is:

-promoted with Rhodium to increase the surface area/area of contact.

 

-added/coated on the surface of asbestos to form platinized asbestos to

reduce the amount/quantity used.

 

The catalyst does not increase the yield of Nitrogen (II)Oxide but it speed up its rate of formation.

 

 

 

C.SAMPLE REVISION QUESTIONS

 

1.State two distinctive features  of a dynamic equilibrium.

 

(i)the rate of forward reaction is equal to the rate of forward reaction

(ii)at equilibrium the concentrations of reactants and products do not change.

 

  1. Explain the effect of increase in pressure on the following:

 

 

(i) N2(g)

 

 

+

 

 

O2(g)  =====  2NO(g)

 

 

 

Gay Lussacs law 1Volume 1Volume 2 Volume Avogadros law 1 molecule 1 molecule 2 molecule

 

 

2 volume on reactant side produce 2 volume on product side.

 

Increase in pressure thus have no effect on position of equilibrium.

 

(ii) 2H2(g)   +    CO(g)    =====     CH3OH (g)

 

Gay Lussacs law             2Volume              1Volume                       1 Volume

Avogadros law                  2 molecule         1 molecule                 1 molecule

3 volume on reactant side produce 1 volume on product side.

 

Increase in pressure shift the equilibrium forward  to the left. More yield of

 

CH3OH is formed.

 

  1. Explain the effect of increasing temperature on the following:

2SO2(g)      +        O2 (g)   =====     2SO3 (g) ΔH = -189kJ

 

Forward reaction is exothermic. Increase in temperature shift the equilibrium backward to reduce the excess heat.

 

5.120g of brass an alloy of copper and Zinc was put it a flask containing dilute hydrochloric acid. The flask was placed on an electric balance. The readings on the balance were recorded as in the table below

 

Time(Seconds) Mass of flask(grams) Loss in mass(grams)
0 600
20 599.50
40 599.12
60 598.84
80 598.66
100 598.54
120 598.50
140 598.50
160 598.50

 

(a)Complete the table by calculating the loss in mass

 

(b)What does the 600 gram reading on the balance represent

 

The initial mass of brass and the acid before any reaction take place.

 

(c)Plot a graph of Time (x-axes) against loss in mass.

 

(d)Explain the shape of your graph

 

The reaction produce hydrogen gas as one of the products that escape to the atmosphere. This decreases the mass of flask.After 120 seconds,the

 

react is complete. No more hydrogen is evolved.The mass of flask remain constant.

 

(d)At what time was the loss in mass equal to:

(i)1.20g

Reading from a correctly plotted graph =

 

(ii)1.30g

Reading from a correctly plotted graph =

 

(iii)1.40g

 

Reading from a correctly plotted graph = (e)What was the loss in mass at:

(i)50oC

Reading from a correctly plotted graph =

 

(ii) 70oC

Reading from a correctly plotted graph =

 

(iii) 90oC g

Reading from a correctly plotted graph =

 

 

21.0.0 ELECTROCHEMISTRY

 

(25 LESSONS)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Electrochemistry can be defined as the study of the effects of electricity on a substance/ compound and how chemical reactions produce electricity. Electrochemistry therefore deals mainly with:

 

  1. Reduction and oxidation
  2. Electrochemical (voltaic) cell
  • Electrolysis (electrolytic) cell

 

(i)REDUCTION AND OXIDATION (REDOX)

 

 

  1. In teams of oxygen transfer:

 

  1. Reduction is removal of oxygen.
  2. Oxidation is addition of oxygen.

 

  • Redox is simultaneous addition and removal of oxygen.

 

  1. Reducing agent is the species that undergoes oxidation, therefore gains

 

  1. Oxidizing agent is the species that undergoes reduction, therefore looses/donates

 

e.g. When hydrogen is passed through heated copper (II) oxide, it is oxidised
to copper metal as in the equation below:
CuO (s) + H2 (g) -> Cu (s) +H2O (l)
1

 

(Oxidising agent)       (Reducing agent)

 

  1. In terms of hydrogen transfer:
    1. Oxidation is the removal of hydrogen.

 

  1. Reduction is the addition of hydrogen.
  • Redox is simultaneous addition and removal of hydrogen.

 

  1. Reducing agent is the species that undergoes oxidation, therefore looses/ donates

 

  1. Oxidizing agent is the species that undergoes reduction, therefore gains

 

e.g. When hydrogen sulphide gas is bubbled into a gas jar containing chlorine gas it is oxidized (loose the hydrogen) to sulphur (yellow solid). The chlorine is reduced (gain hydrogen) to hydrogen chlorine gas.

Cl2 (g) + H2S (g)-> S(S) + 2HCl (g)
(Oxidizing agent) (Reducing agent)

 

  1. In terms of electron transfer:
    1. Oxidation is donation/ loss/ removal of electrons.
    2. Reduction is gain/ accept/ addition of electrons.

 

  • Redox is simultaneous gain/ accept/ addition and donation/ loss/ removal of electrons.

 

  1. Reducing agent is the species that undergoes oxidation, therefore looses/ donates

 

  1. Oxidizing agent is the species that undergoes reduction, therefore gains/ accepts

 

Example

  1. a) Displacement of metals from their solutions:

 

Place 5cm3 each of Iron (II) sulphate (VI) solution into three different test tubes. Add about 1g of copper tunings / powder into one test tube then zinc and magnesium powders separately into the other test tubes. Shake thoroughly for 2 minutes each. Record any colour changes in the table below.

 

Metal added to Iron (II) sulphate (VI) Colour changes
solution
Copper Solution remains green
Zinc Green colour fades
Magnesium Green colour fades
Explanation

 

-When a more reactive metal is added to a solution of less reactive metal, it displaces it from its solution.

 

2

 

-When a less reactive metal is added to a solution of a more reactive metal, it does not displace it from its solution.

 

-Copper is less reactive than iron therefore cannot displace iron its solution.

-Zinc is more reactive than iron therefore can displace iron from its solution.

 

-Magnesium is more reactive than iron therefore can displace iron from its solution.

 

In terms of electron transfer:

 

  • the more reactive metal undergoes oxidation (reducing agent) by donating/loosing electrons to form ions

 

-the less reactive metal undergoes reduction (oxidizing agent) by its ions in solution gaining /accepting/acquiring the electrons to form the metal.

 

-displacement of metals involves therefore electron transfer from a more reactive metal to ions of another less reactive metal.

Examples

 

1. Zn(s)  -> Zn2+(aq) + 2e (oxidation/donation of electrons)
Fe2+(aq) + 2e -> Fe(s) (reduction/gain of electrons)
Fe2+(aq)  + Zn(s) -> Zn2+(aq) +  Fe(s)  (redox/both donation and gain of
electrons)
2. Mg(s) -> Mg2+(aq) + 2e (oxidation/donation of electrons)
Fe2+(aq) + 2e -> Fe(s) (reduction/gain of electrons)
Fe2+(aq)  + Mg(s) -> Mg2+(aq) +  Fe(s)  (redox/both donation and gain of
electrons)
3. Zn(s)  -> Zn2+(aq) + 2e (oxidation/donation of electrons)
Cu2+(aq) + 2e -> Cu(s) (reduction/gain of electrons)
Cu2+(aq)  + Zn(s) -> Zn2+(aq) +  Cu(s)  (redox/both donation and gain of
electrons)
4. Fe(s)  -> Fe2+(aq) + 2e (oxidation/donation of electrons)
2Ag+(aq) + 2e -> 2Ag(s) (reduction/gain of electrons)
2Ag+(aq)  + Fe(s) -> Fe2+(aq) +  2Ag(s)  (redox/both donation and gain of
electrons)
5. Zn(s) -> Zn2+(aq) + 2e (oxidation/donation of electrons)
Cl 2 (g)  + 2e -> 2Cl (aq) (reduction/gain of electrons)
Zn2+(aq) +  2Cl(aq)  (redox/both donation and
Cl2(g)  + Zn(s) -> gain of
electrons)
6. 2Mg(s) -> 2Mg2+(aq) + 4e (oxidation/donation of electrons)
O 2 (g)  + 4e -> 2O 2- (aq) (reduction/gain of electrons)
2Mg2+(aq) +  2O2-(aq)  (redox/both donation and  gain
O2(g)  + 2Mg(s) ->

of electrons)

 

 

3

 

Note

 

(i)The number of electrons donated/lost MUST be equal to the number of electrons gained/acquired.

 

(i)During displacement reaction, the colour of ions /salts fades but does not if displacement does not take place. e.g

 

a)Green colour of Fe2+(aq) fades if Fe2+(aq) ions are displaced from their solution. Green colour of Fe2+(aq) appear if Fe/iron displaces another salt/ions from their solution.

 

b)Blue colour of Cu2+(aq) fades if Cu2+(aq) ions are displaced from their solution and brown copper deposits appear. Blue colour of Cu2+(aq) appear if Cu/copper displaces another salt/ions from their solution.

 

c)Brown colour of Fe3+(aq) fades if Fe3+(aq) ions are displaced from their solution. Brown colour of Fe3+(aq) appear if Fe/iron displaces another salt/ions from their solution to form Fe3+(aq).

 

(iii)Displacement reactions also produce energy/heat. The closer/nearer the metals in the reactivity/electrochemical series the less energy/heat of displacement.

 

(iv)The higher the metal in the reactivity series therefore the easier to loose/donate electrons and thus the stronger the reducing agent.

 

  1. (a)In terms of oxidation number:
    1. Oxidation is increase in oxidation numbers.
    2. Reduction is decrease in oxidation numbers.

 

  • Redox is simultaneous increase in oxidation numbers of one species/substance and a decrease in oxidation numbers of another species/substance.

 

  1. Reducing agent is the species that undergoes oxidation, therefore increases its oxidation number.

 

  1. Oxidizing agent is the species that undergoes reduction, therefore increases its oxidation number.

 

(b)The idea/concept of oxidation numbers uses/applies the following simple guideline rules:

 

Guidelines /rules applied in assigning oxidation number

 

1.Oxidation number of combined Oxygen is always -2 except in peroxides (Na2O2/H2O2) where its Oxidation number is -1

 

 

 

4

 

2.Oxidation number of combined Hydrogen is always +1except in Hydrides (NaH/KH) where its Oxidation number is -1

 

3.All atoms and molecules of elements have oxidation number 0 (zero)

 

Atom Oxidation number Molecule Oxidation number
Na 0 Cl2 0
O 0 O2 0
H 0 H2 0
Al 0 N2 0
Ne 0 O3 0
K 0 P3 0
Cu 0 S8 0

 

4.All combined metals and non-metals have oxidation numbers equal to their valency /oxidation state e.g.

 

Metal/non-metal ion Valency Oxidation state Oxidation number
Fe2+ 2 -2 -2
Fe3+ 3 -3 -3
Cu2+ 2 -2 -2
Cu+ 1 +1 +1
Cl 1 -1 -1
O2- 2 -2 -2
Na+ 1 +1 +1
Al3+ 3 +3 +3
P3- 3 -3 -3
Pb2+ 2 +2 +2

 

5.Sum of oxidation numbers of atoms of elements making a compound is equal zero(0) e.g.

Using this rule ,an unknown oxidation number of an atom in a compound can be

determined as below:

  1. a) CuSO4 has-

-one atom of Cu with oxidation number  +2( refer to Rule 4)

 

-one atom of  S with oxidation number  +6 ( refer to Rule 4)

-six atoms of O each with oxidation number  -2( refer to Rule 4)

Sum of oxidation numbers of atoms in CuSO4 = (+2 + +6 + (-2 x 6)) = 0

 

  1. b) H2SO4 has-

-two atom of H each with oxidation number  +1( refer to Rule 2)

-one atom of  S with oxidation number  +6 ( refer to Rule 4)

-four atoms of O each with oxidation number  -2( refer to Rule 4)

Sum of oxidation numbers of atoms in H2SO4 = (+2 + +6 + (-2 x 4)) = 0

 

5

 

  1. c) KMnO4 has-

-one atom of K with oxidation number  +1( refer to Rule 4)

 

-one atom of  Mn with oxidation number  +7 ( refer to Rule 4)

 

-four atoms of O each with oxidation number -2( refer to Rule 4) Sum of oxidation numbers of atoms in KMnO4 = (+1 + +7 + (-2 x 4)) = 0

 

Determine the oxidation number of:

I.Nitrogen in;

-NO    => x + -2 = 0 thus  x = 0 (-2) = + 2

 

The chemical name of this compound is thus Nitrogen(II)oxide -NO2 => x + (-2 x2)= 0 thus x = 0 (-4) = + 4

 

The chemical name of this compound is thus Nitrogen(IV)oxide -N2O => 2x + -2 = 0 thus 2x = 0 (-2) = +2/2= +1

The chemical name of this compound is thus Nitrogen(I)oxide

 

  1. Sulphur in;

-SO2    => x + (-2 x2)= 0 thus  x = 0 (-4) = + 4

 

The chemical name of this compound is thus Sulphur(IV)oxide -SO3 => x + (-2 x3)= 0 thus x = 0 (-6) = + 6

 

The chemical name of this compound is thus Sulphur(VI)oxide -H2SO4 = ((+1 x 2) + x + (-2 x 4)) thus x= 0-( +2 +-8) =+6 The chemical name of this compound is thus Sulphuric(VI)acid -H2SO3 = ((+1 x 2) + x + (-2 x 3)) thus x= 0-( +2 +-6) =+4 The chemical name of this compound is thus Sulphuric(IV)acid

 

III. Carbon in;

-CO2   => x + (-2 x2)= 0 thus  x = 0 (-4) = + 4

 

The chemical name of this compound is thus carbon(IV)oxide -CO => x + -2 = 0 thus x = 0 -2 = + 2

 

The chemical name of this compound is thus carbon(II)oxide -H2CO3 = ((+1 x 2) + x + (-2 x 3)) thus x= 0-( +2 +-6) =+4 The chemical name of this compound is thus Carbonic(IV)acid

IV.Manganese in;

-MnO2   => x + (-2 x2)= 0 thus  x = 0 (-4) = + 4

 

The chemical name of this compound is thus Manganese(IV)oxide -KMnO4 = ((+1 + x + (-2 x 4)) thus x= 0-( +1 +-8) =+7

 

The chemical name of this compound is thus Potassium manganate(VII) V.Chromium in;

– Cr2O3   => 2x + (-2 x 3)= 0 thus  2x = 0 (-6) = +6 / 2= +3

 

The chemical name of this compound is thus Chromium(III)oxide -K2Cr2O7 => (+1 x 2) + 2x + (-2 x7)= 0

 

thus 2x = 0 +2 +-14 = +12 / 2= +6

 

The chemical name of this compound is thus Potassium dichromate(VI) -K2CrO4 => (+1 x 2) + x + (-2 x4)= 0

 

thus 2x = 0 +2 +-8 = +12 / 2= +6

 

6

 

The chemical name of this compound is thus Potassium chromate(VI)

 

6.The sum of the oxidation numbers of atoms of elements making a charged radical/complex ion is equal to its charge.

 

Using this rule ,the oxidation number of unknown atom of an element in a charged radical/complex ion can be determined as in the examples below;

  1. a) SO42- has-

-one atom of S with oxidation number  +6( refer to Rule 4)

-four atoms of O each with oxidation number  -2( refer to Rule 1)

Sum of oxidation numbers of atoms in SO4 2- = ( +6 + (-2 x 4)) = -2
The chemical name of this radical is thus sulphate(VI) ion
b) NO3 has-
-one atom of N with oxidation number  +4( refer to Rule 4)
-three atoms of O each with oxidation number  -2( refer to Rule 1)
Sum of oxidation numbers of atoms in NO3 = ( +4 + (-2 x 3)) = -1
The chemical name of this radical is thus nitrate(IV) ion.
Determine the oxidation number of:
I.Nitrogen in; => x + (-2 x2)= -1 thus  x = -1 (-4) = + 3
-NO2
The chemical name of this compound/ion/radical is thus Nitrate(III)ion
II. Sulphur in;
-SO3 2- => x + (-2 x3)= -2 thus  x = -2 (-6) = + 4
The chemical name of this compound/ion/radical is thus Sulphate(IV)ion
III. Carbon in;
-CO3 2- = x + (-2 x 3) = -2 thus  x = -2 (-6) = + 4
The chemical name of this compound/ion/radical is thus Carbonate(IV)ion
IV.Manganese in;
-MnO4 =  x + (-2 x 4)= -1 thus x= -1-(-2 +-8) =+7
The chemical name of this compound/ion/radical is thus manganate(VII) ion
V.Chromium in
-Cr2O72-   => 2x + (-2 x7)= -2
thus 2x = -2 +2 +-14 = +12 / 2= +6
The chemical name of this compound/ion//radical is thus dichromate(VI) ion
-CrO4 2- => x + (-2 x4)= -2
thus x =  -2 + (-2 x 4) = +6

The chemical name of this compound/ion//radical is thus chromate(VI) ion

 

(c)Using the concept/idea of oxidation numbers as increase and decrease in oxidation numbers , the oxidizing and reducing species/agents can be determined as in the following examples;

 

(i) Cu2+ (aq) + Zn(s) -> Zn2+ (aq) + Cu(s) Oxidation numbers -> +2 0 +2 0

 

Oxidizing species/agents =>Cu2+ ;its oxidation number decrease from+2 to 0 in Cu(s)

 

7

 

Reducing species/agents => Zn2+ ;its oxidation number increase from 0 to +2 in Zn(s)

 

 

(ii)                                                     2Br   (aq)   +         Cl2(g)  ->          2Cl

 

Oxidation numbers ->        -1                                        0                           -1

 

 

(aq)

 

 

+

 

 

Br2 (l)

0

 

 

Oxidizing agent =>Cl2(g) ;its oxidation number decrease from 0 to-1 in 2Cl                                                                                                                                                                                     (aq)

 

 

(iii)                                                   Br2 (l)              +         Zn(s)  ->          Zn2+   (aq)   +        2Br (aq)

 

Oxidation numbers ->          0                                        0                           +2                                       -1

Oxidizing agent => Br2 (l) ;its oxidation number decrease from 0 to-1 in 2Br(aq) Reducing agents => Zn(s) ;its oxidation number increase from 0 to +2 in Zn2+

 

(iv) 2HCl (aq) + Mg(s) -> MgCl2 (aq) +  H2 (g)
Oxidation numbers -> 2 (+1  -1) 0 +2  2(-1) 0
Oxidizing agent => H+ in HCl;its oxidation number decrease from +1to 0 in H2 (g)
Reducing agents => Mg(s) ;its oxidation number increase from 0 to +2 in Mg2+
(v) 2H2O (l)  + 2Na(s)  -> 2NaOH  (aq) + H2 (g)
Oxidation numbers -> +1  -2 0 +1 -2  +1 0
Oxidizing agent => H+ in H2O;its oxidation number decrease from +1to 0 in H2 (g)
Reducing agents => Na(s) ;its oxidation number increase from 0 to +1 in Na+
(vi)  5Fe2+ (aq) +  8H+ (aq)  +  MnO4 ->  5Fe3+ (aq) +  Mn2+ (aq)  +  4H2O (l)
+2 +1 +7 -2 +3 +2 +1  -2

Oxidizing agent => Mn in MnO4 ;its oxidation number decrease from +7to+2 in Mn2+

Reducing agents => Fe2+ ;its oxidation number increase from +2 to +3 in Fe3+
(vii) 6Fe2+ (aq) + 14H+ (aq) + Cr2O7 2-(aq) -> 6Fe3+ (aq)  +  Cr3+ (aq)  +  7H2O (l)
+2 +1 +6  -2 +3 +3 +1 -2
Oxidizing agent:
Cr in Cr2O72- ;its oxidation number decrease from +6 to+3 in Cr3+
Reducing agents => Fe2+ ;its oxidation number increase from +2 to +3 in Fe3+
(viii) 2Fe2+ (aq)  + 2H+ (aq) +  H2O2(aq) ->  2Fe3+ (aq)  + 2H2O (l)
+2 +1 +1  -1 +3 +1  -2
Oxidizing agent:
O in H2O2;its oxidation number decrease from -1 to -2 in H2O
Reducing agents => Fe2+ ;its oxidation number increase from +2 to +3 in Fe3+
(ix) Cr2O7 2-(aq)  + 6H+ (aq) +  5H2O2(aq) ->  2Cr3+ (aq)  + 2H2O (l) + 5O2(g)
+6  -2 +1 +1  -1 +3 +1  -2 0

Oxidizing agents:

O in H2O2;its oxidation number decrease from -1 to -2 in H2O

Cr in Cr2O72- its oxidation number decrease from +6 to +3 in Cr3+ Reducing agents

 

8

 

O in H2O2;its oxidation number increase from -1 to O in O2(g)

  • in Cr2O72- its oxidation number increase from -2 to O in O2(g)

 

(x) 2MnO (aq)  + 6H+ (aq) + 5HO (aq) ->  2Mn2+ (aq)  + 8H 2 O (l) + 5O (g)
4 2 2 2
+7  -2 +1 +1  -1 +2 +1 -20
Oxidizing agents:

O in H2O2;its oxidation number decrease from -1 to -2 in H2O

 

Mn in MnO4 its oxidation number decrease from +7 to +2 in Mn2+ Reducing agents

O in H2O2;its oxidation number increase from -1 to O in O2(g)

O in MnO4 its oxidation number increase from -2 to O in O2(g)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(ii)ELECTROCHEMICAL (VOLTAIC) CELL

 

 

 

 

9

 

  1. When a metal rod/plate is put in a solution of its own salt, some of the metal ionizes and dissolve into the solution i.e.
M(s) -> M+(aq) + e ( monovalent metal)
M(s) -> M2+(aq) + 2e ( divalent metal)
M(s) -> M3+(aq) + 3e ( Trivalent metal)

 

The ions move into the solution leaving electrons on the surface of the metal rod/plate.

 

2.The metal rod becomes therefore negatively charged while its own solution positively charged. As the positive charges of the solution increase, some of them recombine with the electrons to form back the metal atoms

M+(aq) + e -> M(s) ( monovalent metal)
M2+(aq) + 2e -> M(s) (divalent metal)
M3+(aq) + 3e -> M(s) (Trivalent metal)

 

  1. When a metal rod/plate is put in a solution of its own salt, it constitutes/forms a half-cell. The tendency of metals to ionize differ from one metal to the other. The difference can be measured by connecting two half cells to form an electrochemical/voltaic cell as in the below procedure:

 

To set up an electrochemical /voltaic cell

To compare the relative tendency of metals to ionize

 

Place 50cm3 of 1M Zinc(II) sulphate(VI) in 100cm3 beaker. Put a clean zinc rod/plate into the solution. Place 50cm3 of 1M Copper(II) sulphate(VI) in another 100cm3 beaker. Put a clean copper rod/plate of equal area (length x width) with Zinc into the solution. Connect/join the two metals(to a voltmeter) using connecting wires. Dip a folded filter paper into a solution of Potassium nitrate(V) or sodium(I) chloride(I) until it soaks. Use the folded soaked filter paper to connect/join the two solutions in the two beakers.The whole set up should be as below

 

 

 

 

V

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Repeat the above procedure by replacing:

 

(i)Zinc half cell with Magnesium rod/plate/ribbon dipped in 50cm3 of IM magnesium (II) sulphate(VI) solution

 

10

 

(ii)Zinc half cell with Silver rod/plate/coin dipped in 50cm3 of IM silver(I) nitrate(V) solution

 

(iii)Copper half cell with Iron rod/plate/spoon dipped in 50cm3 of IM Iron (II) sulphate(VI) solution

 

Record the observations in the table below

Changes on the Changes on Changes on the Changes on the Voltage/voltme
1st  metal rod the  2nd   metal 1st  solution 2nd solution ter
(A) rod (B) (A(aq)) (B(aq)) reading(Volts)
Using Zn/Cu
half cell
-The rod -copper rod Zinc(II)sulphate Blue Copper 0.8
decrease in size /plate increase (VI)colour (II)sulphate (Theoretical
/mass in size /mass/ remain (VI)colour value=1.10V)
/dissolves/ deposited colourless fades. Brown
erodes solid/residue/
deposit
Using Mg/Cu
half cell
-The rod -copper rod Magnesium(II) Blue Copper 1.5
decrease in size /plate increase sulphate(VI) (II)sulphate (Theoretical
/mass in size /mass/ colour remain (VI)colour value=2.04V)
/dissolves/ deposited colourless fades Brown
erodes solid/residue/
deposit
Using Ag/Cu
half cell
-The rod -silver coin/ Blue Copper Silver(I)nitrate 0.20
increase in size rod /plate (II)sulphate (V)colour (Theoretical
/mass /deposited increase in (VI)colour remain value=0.46V)
size /mass/ remains colourless
deposited
Using Fe/Cu
half cell
-The rod -copper rod Iron(II)sulphate Blue Copper 0.60
decrease in size /plate increase (VI)colour (II)sulphate (Theoretical
/mass in size /mass/ becomes more (VI)colour value=0.78V)
/dissolves/ deposited green fades.Brown
erodes solid/residue/
deposit

From the above observations ,it can be deduced that:

 

(i)in the Zn/Cu half-cell the;

-Zinc rod/plate ionizes /dissolves faster than the copper rod/plate to form Zn2+

11

 

Ionic equation Zn(s) -> Zn2+(aq) + 2e
-blue copper ions in the Copper (II)sulphate solution gains the donated
electrons to form brown copper metal/atoms
Ionic equation Cu2+(aq) + 2e-> Cu(s)

 

This reaction shows /imply the Zinc rod has a higher tendency to ionize than copper.The Zinc rod has a higher net accumulation of electrons and is more negative compared to the copper rod which has lower accumulation of electrons. The copper rod is therefore relatively more positive with respect to Zinc rod.

 

When the two half cells are connected , electrons therefore flow from the negative Zinc rod through the external wire to be gained by copper ions. This means a net accumulation/increase of Zn2+ positive ions on the negative half cell and a net decrease in Cu2+ positive ions on the positive half cell.

The purpose of the salt bridge therefore is:

 

(i)complete the circuit

(ii)maintain balance of charges /ions on both half cells.

For the negative half cell the NO3 /Cl from salt bridge decrease/neutralise the increased positive(Zn2+) ion.

 

For the positive half cell the Na+ / K+ from salt bridge increase the decreased positive(Cu2+) ion.

 

The voltmeter should theoretically register/read a 1.10Volts as a measure of the electromotive force (e.m.f) of the cell .Practically the voltage reading is lowered because the connecting wires have some resistance to be overcomed.

 

A combination of two half cells that can generate an electric current from a redox reaction is called a voltaic/electrochemical cell.

 

By convention a voltaic/electrochemical cell is represented;

M(s)         /  M2+(aq)  //   N2+ (aq)  /    N(s)

(metal rod of M)(solution ofM)(solution ofN)(metal rod ofN)

 

Note;

a)(i)Metal M must be the one higher in the reactivity series.

(ii)It forms the negative terminal of the cell.

 

(iii)It must diagrammatically be drawn first on the left hand side when illustrating the voltaic/electrochemical cell.

 

b)(i)Metal N must be the one lower in the reactivity series.

(ii)It forms the positive terminal of the cell.

 

(iii)It must diagrammatically be drawn second/after/ right hand side when illustrating the voltaic/electrochemical cell.

 

Illustration of  the voltaic/electrochemical cell.

 

(i)Zn/Cu cell

 

12

 

  1. Zinc rod ionizes /dissolves to form Zn2+ ions at the negative terminal

Zn(s)               ->            Zn2+(aq)     +              2e

 

  1. Copper ions in solution gain the donated electrons to form copper atoms/metal

Cu2+(aq)+2e->Cu(s)

 

3.Overall redox equation

 

 

Cu2+(aq)

 

 

+    Zn(s)

 

->

 

Zn2+(aq)

 

+

 

 

Cu(s)

 

 

4.cell representation.

 

 

Zn(s) / 1M, Zn2+

 

 

(aq)

 

// 1M,Cu2+(aq) / Cu(s)  E0

 

 

= +1.10 V

 

 

5.cell  diagram

 

 

Voltmeter

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(ii)Mg/Cu cell

 

  1. Magnesium rod ionizes /dissolves to form Mg2+ ions at the negative terminal

Mg(s)              ->            Mg2+(aq)   +              2e

 

  1. Copper ions in solution gain the donated electrons to form copper atoms/metal

Cu2+(aq)+2e->Cu(s)

 

3.Overall redox equation

Cu2+(aq)     +    Mg(s)        ->         Mg2+(aq)   +            Cu(s)

 

4.cell representation.

Mg(s) / 1M, Mg2+(aq) // 1M,Cu2+(aq) / Cu(s)  E0  = +2.04 V

 

5.cell  diagram.

 

 

 

 

 

13

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(iii)Fe/Cu cell

 

  1. Magnesium rod ionizes /dissolves to form Mg2+ ions at the negative terminal

Fe(s)                ->            Fe2+(aq)     +              2e

 

  1. Copper ions in solution gain the donated electrons to form copper atoms/metal

Cu2+(aq)+2e->Cu(s)

 

3.Overall redox equation

Cu2+(aq)     +    Fe(s)       ->           Fe2+(aq)     +            Cu(s)

 

4.cell representation.

Fe(s) / 1M, Fe2+(aq) // 1M,Cu2+(aq) / Cu(s)  E0  = +0.78 V

 

5.cell  diagram.

 

V

 

Fe

 

 

 

 

 

 

Fe++

 

 

 

 

 

 

14

 

(iv)Ag/Cu cell

 

  1. Copper rod ionizes /dissolves to form Cu2+ ions at the negative terminal

Cu(s)              ->            Cu2+(aq)    +              2e

 

  1. Silver ions in solution gain the donated electrons to form silver atoms/metal

2Ag+(aq)    +2e->2Ag(s)

 

3.Overall redox equation

2Ag+(aq)     +    Cu(s)       ->          Cu2+(aq)    +            2Ag(s)

 

4.cell representation.

Cu(s) / 1M, Cu2+(aq) // 1M,2Ag+(aq) / 2Ag(s)  E0  = +0.46 V

 

5.cell  diagram.

 

V
Cu Ag

 

 

 

 

 

 

Ag+

Cu++

 

 

 

Cu(s)  +  2Ag+ (aq)            Cu2+(aq) + 2Ag(s)

 

Standard electrode potential  (Eᶿ)

 

The standard electrode potential (Eᶿ) is obtained if the hydrogen half cell is used as reference. The standard electrode potential (Eᶿ) consist of inert platinum electrode immersed/dipped in 1M solution of (sulphuric(VI) acid) H+ ions. Hydrogen gas is bubbled on the platinum electrodes at:

(i)a temperature of 25oC

 

(ii)atmospheric pressure of 101300Pa/101300Nm-2/1atm/760mmHg/76cmHg (iii)a concentration of 1M(1moledm-3) of sulphuric(VI) acid/ H+ ions and

1M(1moledm-3) of the other half cell.

 

Hydrogen is adsorbed onto the surface of the platinum. An equilibrium/balance exist between the adsorbed layer of molecular hydrogen and H+ ions in solution to

form a half cell.

 

 

½ H2 (g)

 

====

 

H+ (aq)

 

+

 

e

 

 

 

The half cell representation is:

Pt,½ H2 (g)  / H+ (aq), 1M

 

 

15

 

The standard electrode potential (Eᶿ) is thus defined as the potential difference for a cell comprising of a particular element in contact with1M solution of its own ions and the standard hydrogen electrode.

 

If the other electrode has a higher/greater tendency to lose electrons than the hydrogen electrode, the electrode is therefore negative with respect to hydrogen electrode and its electrode potential has negative (Eᶿ) values.

 

If the other electrode has a lower/lesser tendency to lose electrons than the hydrogen electrode, the electrode is therefore positive with respect to hydrogen electrode and its electrode potential has positive (Eᶿ) values.

 

Table showing the standard electrode potential (Eᶿ) of some reactions

 

Reaction (Eᶿ) values in volts
F2 (g)+ 2e -> 2F (aq) +2.87
H2 O2 (aq)+ H+ (aq) +2e -> H2 O (l) +1.77
Mn O4 (aq)+ 4H+ (aq) +3e -> MnO2 (s) +H2 O (l) +1.70
2HClO (aq)+ 2H+ (aq) +2e -> Cl2 (aq) +2H2 O (l) +1.59
Mn O4 (aq)+ 4H+ (aq) +5e -> Mn2+ (aq) +H2 O (l) +1.51
Cl2 (g)+ 2e -> 2Cl (aq) +1.36
Mn O2 (s)+ 4H+ (aq) +2e -> Mn2+ (aq) +2H2 O (l) +1.23
Br2 (aq)+ 2e -> 2Br (aq) +1.09
NO3 (aq)+ 2H+ (aq) +  e -> NO2 (g) + H2 O (l) +0.80
Ag+ (aq)  +  e ->  Ag(s) +0.80
Fe3+ (aq) + e ->  Fe2+ (aq) +0.77
2H+ (aq)+ O2 (g)  -> H2 O2 (aq) +0.68
I2 (aq)+ 2e -> 2I (aq) +0.54
Cu2+ (aq)  +  2e ->  Cu(s) +0.34
2H+ (aq)  +  2e ->  H2(g) +0.00
Pb2+ (aq) + 2e -> Pb(s) -0.13
Fe2+ (aq) + 2e -> Fe(s) -0.44
Zn2+ (aq) + 2e ->  Zn(s) -0.77
Al3+ (aq) + 3e -> Al(s) -1.66
Mg2+ (aq)  +  2e ->  Mg(s) -2.37
Na+ (aq) + e -> Na(s) -2.71
K+ (aq) + e -> K(s) -2.92

 

Note:

(i)Eᶿ values generally show the possibility/feasibility of a reduction process/oxidizing strength.

 

(ii)The element/species in the half cell with the highest negative Eᶿ value easily gain / acquire electrons.

 

It is thus the strongest oxidizing agent and its reduction process is highly possible/feasible. The element/species in the half cell with the lowest positive Eᶿ value easily donate / lose electrons.

 

16

 

It is thus the strongest reducing agent and its reduction process is the least possible/feasible.

 

(iii)The overall redox reaction is possible/feasible is it has a positive (+) Eᶿ. If the overall redox reaction is not possible/ not feasible/ forced, it has a

 

negative (-) Eᶿ

 

Calculation examples on Eᶿ

 

Calculate the Eᶿ value of a cell made of:

a)Zn and Cu

From the table above:

 

Cu2+ (aq) + 2e -> Cu(s) Eᶿ = +0.34V(higher Eᶿ /Right Hand Side diagram) Zn2+ (aq) + 2e ->Zn(s) Eᶿ = -0.77V(lower Eᶿ/ Left Hand Side diagram) Zn(s) ->Zn2+ (aq) + 2e Eᶿ = +0.77(reverse lower Eᶿ to derive cell reaction / representation)

 

Overall Eᶿ = Eᶿ higher- Eᶿ lower / Eᶿ RHS – Eᶿ LHS/ Eᶿoxidized- Eᶿ reduced

 

Substituting:

Overall Eᶿ = +0.34 (- 0.77) = +1.10V
Overall redox equation:
Cu2+ (aq) + Zn(s)  ->  Zn2+ (aq)  +  Cu(s)   Eᶿ = +1.10V
Overall conventional cell representation: Eᶿ  = +1.10V
Zn(s) / Zn2+ (aq) 1M,  // 1M,Cu2+ (aq)  / Cu(s)
Overall conventional cell diagram:
Voltmeter(1.10V)

 

 

 

 

 

 

 

 

 

 

Zn2+

 

1M Cu2+ (aq)

 

1M Zn2+ (aq)

 

 

 

 

 

 

 

 

 

 

 

17

 

Zinc and copper reaction has a positive(+) overall Eᶿ therefore is possible/feasible and thus Zinc can displace/reduce Copper solution.

 

b)Mg and Cu

From the table above:

 

Cu2+ (aq) + 2e -> Cu(s) Eᶿ = +0.34V(higher Eᶿ /Right Hand Side diagram) Mg2+ (aq) + 2e ->Mg(s) Eᶿ = -2.37V(lower Eᶿ/ Left Hand Side diagram) Mg(s) ->Mg2+ (aq) + 2e Eᶿ = +2.37(reverse lower Eᶿ to derive cell reaction / representation)

 

Overall Eᶿ = Eᶿ higher- Eᶿ lower / Eᶿ RHS – Eᶿ LHS/ Eᶿ oxidized- Eᶿ reduced

 

Substituting:

Overall Eᶿ = +0.34 (- 2.37) = +2.71V
Overall redox equation:
Cu2+ (aq) + Mg(s)  ->  Mg2+ (aq)  +  Cu(s)   Eᶿ = +2.71V
Overall conventional cell representation: Eᶿ  = +2.71V
Mg(s) / Mg2+ (aq) 1M,  // 1M,Cu2+ (aq)  / Cu(s)

 

c)Ag and Pb

From the table above:

2Ag+ (aq) + 2e -> 2Ag(s) Eᶿ = +0.80V(higher Eᶿ /Right Hand Side diagram) Pb2+ (aq) + 2e ->Pb(s) Eᶿ = -0.13V(lower Eᶿ/ Left Hand Side diagram) Pb(s) ->Pb2+ (aq) + 2e Eᶿ = +0.13(reverse lower Eᶿ to derive cell reaction / representation)

 

Overall Eᶿ = Eᶿ higher- Eᶿ lower / Eᶿ RHS – Eᶿ LHS/ Eᶿ oxidized- Eᶿ reduced

 

Substituting:

Overall Eᶿ = +0.80 (- 0.13) = +0.93V
Overall redox equation:
2Ag+ (aq) + Pb(s)  ->  Pb2+ (aq)  +  2Ag(s)   Eᶿ = +0.93V
Overall conventional cell representation: Eᶿ  =  +0.93V
Pb(s) / Pb2+ (aq) 1M,  // 1M,2Ag+ (aq)  / Ag(s)

 

d)Chlorine and Bromine

From the table above:

2e  +  Cl2(g) ->2Cl (aq)   Eᶿ = +1.36V(higher Eᶿ /Right Hand Side diagram)

 

2e + Br2(aq) ->2Br (aq) Eᶿ = +0.13V(lower Eᶿ/ Left Hand Side diagram) 2Br (aq) -> Br2(aq) + 2e Eᶿ = -0.13(reverse lower Eᶿ to derive cell reaction / representation)

 

 

Overall Eᶿ = Eᶿ higher- Eᶿ lower / Eᶿ RHS – Eᶿ LHS/ Eᶿ oxidized- Eᶿ reduced

 

Substituting:

Overall Eᶿ =  – 0.13 (- 1.36) = +1.23V

 

Overall redox equation:

 

18

 

2Br (aq) + Cl2(g)  ->  2Cl (aq)  +  Br2(aq)   Eᶿ = +1.23V
Overall conventional cell representation: Eᶿ  =  +1.23V
Cl2(g) / 2Cl (aq) 1M,  // 1M, 2Br (aq)  / Br2(aq)

 

 

Chlorine displaces bromine from bromine water. When chlorine gas is thus bubbled in bromine water, the pale green colour fades as displacement takes place and a brown solution containing dissolved bromine liquid is formed. This reaction is feasible /possible because the overall redox reaction has a positive Eᶿ value.

 

 

 

 

e)Strongest oxidizing agent and the strongest reducing agent.

From the table above: Eᶿ = +2.87V(highest Eᶿ /strongest oxidizing agent)
2e + F2(g) ->2F (aq)
2e + 2K+ (aq) ->2K (aq) Eᶿ = -2.92V(lowest Eᶿ/ strongest reducing agent)
2K (aq) -> 2K+ (aq) + 2e Eᶿ = +2.92V (reverse lower Eᶿ to derive cell reaction /
representation)

 

Overall Eᶿ = Eᶿ higher- Eᶿ lower / Eᶿ RHS – Eᶿ LHS/ Eᶿ oxidized- Eᶿ reduced

 

Substituting:

Overall Eᶿ =  +2.87 (-2.92) = +5.79V
Overall redox equation:
F2(g) + 2K(s)  ->  2F (aq)  +  2K+ (aq)   Eᶿ = +5.79V
Overall conventional cell representation:
2K(s) / 2K+ (aq),1M,  // 1M, 2F (aq)  / F2(g) Eᶿ  = +5.79V

 

The redox reactions in an electrochemical/voltaic is commercially applied to make the:

(a)Dry /primary/Laclanche cell.

 

(b)Wet /secondary /accumulators.

 

(a)Dry/primary/Laclanche cell

 

Examine a used dry cell.

 

Note the positive and the negative terminal of the cell. Carefully using a knife cut a cross section from one terminal to the other.

 

The dry cell consist of a Zinc can containing a graphite rod at the centre surrounded by a paste of;

-Ammonium chloride

-Zinc chloride

-powdered manganese (IV) oxide mixed with Carbon.

 

Zinc acts/serve as the negative terminal where it ionizes/dissociates:

 

19

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Zn(s)  ->       Zn2+(aq)  +  2e

 

Ammonium ions in ammonium chloride serve as the positive terminal where it is converted to ammonia gas and hydrogen gas.

2NH4+(aq)  +  2e  -> 2NH3(g) + H2(g)

 

Ammonia forms a complex salt / compound /(Zn(NH3) 4)2+ (aq) / tetramminezinc(II) complex with the Zinc chloride in the paste.

 

Manganese (IV) oxide oxidizes the hydrogen produced at the electrodes to water preventing any bubbles from coating the carbon terminal which would reduce the efficiency of the cell.

 

Ammonium chloride is used as paste because the solid does not conduct electricity because the ions are fused/not mobile.

 

Since the reactants are used up, the dry /primary /Laclanche cell cannot provide continous supply of electricity.The process of restoring the reactants is called recharging.

 

b)Wet/Secondary/Accumulators

 

  1. Wet/Secondary/Accumulators are rechargeable unlike dry /primary /Laclanche cells.Wet/Secondary/Accumulators are made up of:

 

(i)Lead plate that forms the negative terminal

(ii)Lead(IV) oxide that forms the positive terminal

 

2.The two electrodes are dipped in concentrated sulphuric(VI) acid of a relative density 1.2/1.3

3.At the negative terminal,lead ionizes /dissolves;

Pb(s)      ->  Pb2+   +  2e

 

4.At the positive terminal,

 

20

 

(i) Lead(IV) oxide reacts with the hydrogen ions in sulphuric(VI)acid to form Pb2+ (aq) ions;

PbO2(s)   + 4H+(aq) + 2e ->  Pb2+ (aq)  + H2O(l)

 

(ii) Pb2+ (aq) ions formed instantly react with sulphate (VI) ions/ SO42- (aq) from sulphuric (VI)acid to form insoluble Lead(II) sulphate (VI).

Pb2+ (aq) +  SO42- (aq) ->  PbSO4(s)

 

5.The overall cell reaction is called discharging

PbO2(s)  +Pb(s) + 4H+(aq) + 2SO42- (aq)-> 2PbSO4(s) + 2H2O(l) Eᶿ = +2.0V

 

6.The insoluble Lead(II) sulphate (VI) formed should not be left for long since fine Lead(II) sulphate (VI) will change to a course non-reversible and inactive form making the cell less efficient.

 

As the battery discharges ,lead and lead(IV)oxide are depleted/finished/reduced and the concentration of sulphuric(VI)acid decreases.

 

  1. During recharging, the electrode reaction is reversed as below:

2PbSO4(s) + 2H2O(l) ->PbO2(s)  +Pb(s) + 4H+(aq) + 2SO42- (aq)

 

  1. A car battery has six Lead-acid cells making a total of 12 volts.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(iii)ELECTROLYSIS (ELECTROLYTIC CELL)

 

1.Electrolysis is defined simply as the decomposition of a compound by an electric current/electricity.

 

A compound that is decomposed by an electric current is called an electrolyte. Some electrolytes are weak while others are strong.

 

2.Strong electrolytes are those that are fully ionized/dissociated into (many) ions.

 

Common strong electrolytes include:

(i)all mineral acids

(ii)all strong alkalis/sodium hydroxide/potassium hydroxide.

(iii)all soluble salts

21

 

3.Weak electrolytes are those that are partially/partly ionized/dissociated into (few) ions.

Common weak electrolytes include:

 

(i)all organic acids

(ii)all bases except sodium hydroxide/potassium hydroxide.

(iii)Water

 

  1. 4. A compound that is not decomposed by an electric current is called non-electrolyte. Non-electrolytes are those compounds /substances that exist as molecules and thus cannot ionize/dissociate into(any) ions . Common non-electrolytes include:

 

  • most organic solvents (e.g. petrol/paraffin/benzene/methylbenzene/ethanol) (ii)all hydrocarbons(alkanes /alkenes/alkynes)

 

(iii)Chemicals of life(e.g. proteins, carbohydrates, lipids, starch, sugar)

 

  1. 5. An electrolytes in solid state have fused /joined ions and therefore do not conduct electricity but the ions (cations and anions) are free and mobile in molten and aqueous (solution, dissolved in water) state.

 

6.During electrolysis, the free ions are attracted to the electrodes. An electrode is a rod through which current enter and leave the electrolyte during electrolysis. An electrode that does not influence/alter the products of electrolysis is called an inert electrode.

 

Common inert electrodes include:

(i)Platinum

(ii)Carbon graphite

Platinum is not usually used in a school laboratory because it is very expensive.

Carbon graphite is easily/readily and cheaply available (from used dry cells).

 

7.The positive electrode is called Anode.The anode is the electrode through which current enter the electrolyte/electrons leave the electrolyte

 

8.The negative electrode is called Cathode. The cathode is the electrode through which current leave the electrolyte / electrons enter the electrolyte

 

  1. 9. During the electrolysis, free anions are attracted to the anode where they lose /donate electrons to form neutral atoms/molecules. i.e.

 

M(l)   ->          M+(l)  +  e  (for cations from molten electrolytes)

M(s) -> M+(aq) + e (for cations from electrolytes in aqueous state / solution / dissolved in water)

 

The neutral atoms /molecules form the products of electrolysis at the anode. This is called discharge at anode

 

 

 

22

 

  1. During electrolysis, free cations are attracted to the cathode where they gain

/accept/acquire electrons to form neutral atoms/molecules.

 

X+ (aq) + 2e -> X(s) (for cations from electrolytes in aqueous state / solution / dissolved in water)

2X+ (l)  +  2e -> X (l)  (for cations from molten electrolytes)

 

The neutral atoms /molecules form the products of electrolysis at the cathode. This is called discharge at cathode.

 

  1. The below set up shows an electrolytic cell.

 

 

 

Simple set up of electrolytic cell

 

 

 

Gaseous product at cathode   Gaseous product at anode

 

 

Electrolyte

 

Cathode(-) Anode(+)

 

 

 

 

 

 

Battery

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

23

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

  1. For a compound /salt containing only two ion/binary salt the products of electrolysis in an electrolytic cell can be determined as in the below examples:

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

a)To determine the products of electrolysis of molten Lead(II)chloride

 

(i)Decomposition of electrolyte into free ions;

PbCl2 (l)           ->        Pb 2+(l)  +     2Cl(l)

(Compound decomposed into free cation and anion in liquid state)

 

(ii)At the cathode/negative electrode(-);

Pb 2+(l)   +        2e        ->      Pb (l)

(Cation / Pb 2+ gains / accepts / acquires electrons to form free atom)

 

24

 

(iii)At the anode/positive electrode(+);

2Cl(l)         ->    Cl2 (g)  +   2e

(Anion / Cl donate/lose electrons to form free atom then a gas molecule)

 

(iv)Products of electrolysis therefore are;

I.At the cathode grey beads /solid lead metal.

II.At the anode pale green chlorine gas.

 

 

 

b)To determine the products of electrolysis of molten Zinc bromide

 

(i)Decomposition of electrolyte into free ions;

ZnBr2 (l)           ->       Zn 2+(l)  +     2Br(l)

(Compound decomposed into free cation and anion in liquid state)

 

(ii)At the cathode/negative electrode(-);

Zn 2+(l)   +         2e        ->    Zn(l)

(Cation / Zn2+ gains / accepts / acquires electrons to form free atom)

 

(iii)At the anode/positive electrode(+);

2Br(l)        ->    Br2 (g)  +     2e

 

(Anion / Br donate/lose electrons to form free atom then a liquid molecule which change to gas on heating)

 

(iv)Products of electrolysis therefore are;

I.At the cathode grey beads /solid Zinc metal.

II.At the anode red bromine liquid / red/brown bromine gas.

 

c)To determine the products of electrolysis of molten sodium chloride

 

(i)Decomposition of electrolyte into free ions;

NaCl (l)            ->         Na +(l)  +     Cl(l)

(Compound decomposed into free cation and anion in liquid state)

 

(ii)At the cathode/negative electrode(-);

2Na+(l)   +         2e        ->      Na (l)

(Cation / Na+ gains / accepts / acquires electrons to form free atom)

 

(iii)At the anode/positive electrode(+);

2Cl(l)         ->    Cl2 (g)  +   2e

(Anion / Cl donate/lose electrons to form free atom then a gas molecule)

 

(iv)Products of electrolysis therefore are;

 

I.At the cathode grey beads /solid sodium metal.

II.At the anode pale green chlorine gas.

 

25

 

d)To determine the products of electrolysis of molten Aluminium (III)oxide

 

(i)Decomposition of electrolyte into free ions;

Al2O3 (l)          ->       2Al 3+(l)  +     3O2-(l)

(Compound decomposed into free cation and anion in liquid state)

 

(ii)At the cathode/negative electrode(-);

4Al 3+ (l)   +         12e       ->    4Al (l)

(Cation / Al 3+ gains / accepts / acquires electrons to form free atom)

 

(iii)At the anode/positive electrode(+);

6O2-(l)        ->    3O2 (g)  +     12e

(Anion /6O2- donate/lose 12 electrons to form free atom then three gas molecule)

 

(iv)Products of electrolysis therefore are;

I.At the cathode grey beads /solid aluminium metal.

II.At the anode colourless  gas that relights/rekindles glowing splint.

 

  1. 13. For a compound /salt mixture containing many ions in an electrolytic cell, the discharge of ions in the cell depend on the following factors:

 

  1. Position of cations and anions in the electrochemical series

 

  1. Most electropositive cations require more energy to reduce (gain electrons) and thus not readily discharged. The higher elements /metals in the electrochemical series the less easily/readily it is discharged at the cathode in the electrolytic cell.

 

Table I showing the relative ease of discharge of cations in an electrolytic cell

 

K+(aq) + e -> K(s) (least readily/easily discharged)
Na+(aq) +  e -> Na(s)
Ca2+(aq) +  2e -> Ca(s)
Mg2+(aq) +  2e -> Mg(s)
Al3+(aq) + 3e -> Al(s)
Zn2+(aq) +  2e -> Zn(s)
Fe2+(aq) + 2e -> Fe(s)
Pb2+(aq) + 2e -> Pb(s) (hydrogen is usually metallic)
2H+(aq) + 2e -> H2(g)
Cu2+(aq) +  2e -> Cu(s)
Hg2+(aq) +  2e -> Hg(s)
Ag+(aq) + e -> Ag(s) (most readily/easily discharged)

 

2.The OH ion is the most readily/easily discharged anion . All the other anionic radicals(SO42- ,SO32- ,CO32- ,HSO4 ,HCO3 ,NO3 ,PO43-)are not/never discharged. The ease of discharge of halogen ions increase down the group.

 

Table II showing the relative ease of discharge of anions in an electrolytic cell

 

26

 

4OH (aq) -> 2H2O(l) + O2 (g) + 4e (most readily/easily discharged)
2 I(aq) -> I2(aq) +  2e
2 Br(aq) -> Br2(aq) +  2e
2 Cl(aq) -> Cl2(aq) +  2e
2 F(aq) -> F2(aq) +  2e
SO4 2- ,SO3 2- ,CO32- ,HSO4 ,HCO3 ,NO3 ,PO4 3- not/never/rarely discharged.

 

3.(a)When two or more cations are attracted to the cathode, the ion lower in the electrochemical series is discharged instead of that which is higher as per the table I above. This is called selective/preferential discharge at cathode.

 

(b)When two or more anions are attracted to the anode, the ion higher in the electrochemical series is discharged instead of that which is lower as per the table I above. This is called selective/preferential discharge at anode.

 

4.The following experiments shows the influence /effect of selective/preferential discharge on the products of electrolysis:

 

(i)Electrolysis of acidified water/dilute sulphuric(VI) acid

 

Fill the Hoffmann voltameter with dilute sulphuric(VI) acid. Connect the Hoffmann voltameter to a d.c. electric supply. Note the observations at each electrode.

 

Electrolytic cell set up during electrolysis of acidified water/dilute sulphuric(VI) acid

 

 

 

Simple set up of electrolytic cell

 

 

 

Gaseous product at cathode Gaseous product at anode
Electrolyte
Cathode(-) Anode(+)

 

 

 

 

Battery

 

 

 

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

 

-> OH- (aq)

 

+ H+(aq)

 

 

27

 

H2 SO4(aq)   -> SO42-(aq)  + 2H+(aq)

 

  1. Name the ions in acidified water that are attracted/move to: Cathode- H+(aq) from either sulphuric(VI) acid (H2 SO4) or water (H2O) Anode SO42-(aq) from sulphuric (VI) acid (H2 SO4) and OH (aq) from water

(H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      4H+(aq)           +   4e   ->  2H2(g)

Anode           4OH- (aq)   ->    2H2O(l)   + O2 (g)  +  4e

 

(4OH- ions selectively discharged instead of SO42- ions at the anode)

 

  1. Name the products of electrolysis of acidified water. Cathode-Hydrogen gas (colourless gas that extinguishes burning splint with explosion/ pop sound

AnodeOxygen gas (colourless gas that relights /rekindles glowing splint)

 

  1. Explain the difference in volume of products at the cathode and anode.

 

The four(4) electrons donated/lost by OH ions to form 1 molecule/1volume/1mole of oxygen (O2)gas at the anode are gained/acquired/accepted by the four H+(aq) ions to form 2 molecule/2volume/2mole of Hydrogen (H2)gas at the cathode.

 

The volume of Oxygen gas at the anode is thus a half the volume of Hydrogen produced at the cathode/ The volume of Hydrogen gas at the cathode is thus a twice the volume of Oxygen produced at the anode.

 

  1. Why is electrolysis of dilute sulphuric(VI) acid called electrolysis of (acidified) water?

 

The ratio of H2 (g): O2 (g) is 2:1 as they are combined in water. This implies/means that water in the electrolyte is being decomposed into hydrogen and Oxygen gases. The electrolysis of dilute sulphuric acid is therefore called electrolysis of acidified water.

 

  1. Explain the changes in concentration of the electrolyte during electrolysis of acidified water

 

The concentration of dilute sulphuric (VI) acid increases. Water in the electrolyte is decomposed into Hydrogen and Oxygen gases that escape. The concentration /mole of acid present in a given volume of solution thus continue increasing/rising.

 

(ii)Electrolysis of  Magnesium sulphate(VI) solution

 

Fill the Hoffmann voltameter with dilute sulphuric(VI) acid. Connect the Hoffmann voltameter to a d.c. electric supply. Note the observations at each electrode.

 

 

 

Answer the following questions:

 

 

28

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

Mg SO4(aq)

 

-> OH- (aq)    + H+(aq)

 

-> SO42-(aq)  + Mg2+(aq)

 

 

  1. Name the ions in Magnesium sulphate(VI) solution that are attracted/move

to:

Cathode- Mg2+(aq) from Magnesium sulphate(VI) solution (Mg SO4) and H+(aq) from water (H2O)

 

Anode SO42-(aq) from Magnesium sulphate(VI) solution (Mg SO4) and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      4H+(aq)           +   4e     ->  2H2(g)

H+ ions selectively discharged instead of Mg2+ ions at the cathode)

 

Anode           4OH- (aq)   ->    2H2O(l)   + O2 (g)  +  4e

 

(4OH- ions selectively discharged instead of SO42- ions at the anode)

 

  1. Name the products of electrolysis of Magnesium sulphate(VI) solution

 

Cathode-Hydrogen gas (colourless gas that extinguishes burning splint with explosion/ pop sound

AnodeOxygen gas (colourless gas that relights /rekindles glowing splint)

 

  1. Explain the difference in volume of products at the cathode and anode.

 

The four(4) electrons donated/lost by OH ions to form 1 molecule/1volume/1mole of oxygen (O2)gas at the anode are gained/acquired/accepted by the four H+(aq) ions to form 2 molecule/2volume/2mole of Hydrogen (H2)gas at the cathode.

 

The volume of Oxygen gas at the anode is thus a half the volume of Hydrogen produced at the cathode/ The volume of Hydrogen gas at the cathode is thus a twice the volume of Oxygen produced at the anode.

 

  1. Explain the changes in concentration of the electrolyte during electrolysis of Magnesium sulphate(VI) solution

 

The concentration of dilute Magnesium sulphate(VI) solution increases.

The ratio of H2 (g): O2 (g) is 2:1 as they are combined in water.

 

Water in the electrolyte is decomposed into Hydrogen and Oxygen gases that escape as products.

 

The concentration /mole of acid present in a given volume of Magnesium sulphate(VI) solution thus continue increasing/rising.

 

 

 

 

29

 

 

The set up below was used during the electrolysis of aqueous magnesium sulphate using inert electrodes.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Name a suitable pair of electrodes for this experiment

 

Identify the ions and cations in the solution

 

On the diagram label the cathode

 

Write ionic equations for the reactions that took place at the anode.

 

Explain the change that occurred to the concentration of magnesium sulphate solution during the experience.

 

During the electrolysis a current of 2 amperes was passed through the solution for 4 hours. Calculate the volume of the gas produced at the anode.(1 faraday 96500 coulombs and volume of a gas at room temperature is 24000cm3)

 

One of the uses of electrolysis is electroplating

What is meant by electroplating?

Give tow reasons why electroplating is necessary.

 

  1. Concentration of the electrolytes

 

1.High concentrations of cations and/or anions at the electrodes block the ion/s that is likely to be discharged at the electrode. This is called over voltage. A concentrated solution therefore produces different products of electrolysis from a dilute one.

 

  1. The following experiments show the influence/effect of concentration of electrolyte on the products of electrolysis.

 

(i)Electrolysis of  dilute and concentrated(brine)sodium chloride solution

 

30

 

  1. I. Dissolve about 0.5 g of pure sodium chloride crystals in 100cm3 of water. Place the solution in an electrolytic cell. Note the observations at each electrode for 10 minutes. Transfer the set up into a fume chamber/open and continue to make observations for a further 10 minute.

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

H2O(l)              -> OH- (aq)    + H+(aq)

NaCl(aq)   -> Cl-(aq)  + Na+(aq)

 

  1. Name the ions in sodium chloride solution that are attracted/move to: Cathode- Na+(aq) from Sodium chloride solution (NaCl) and H+(aq) from water

(H2O)

Anode Cl(aq) from sodiumchloride solution (NaCl) and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      4H+(aq)           +   4e     ->  2H2(g)

H+ ions selectively discharged instead of Na+ ions at the cathode)

 

Anode           4OH (aq)   ->    2H2O(l)    + O2 (g)  +  4e

(4OH ions selectively discharged instead of Cl ions at the anode)

 

  1. Name the products of electrolysis of dilute sodium chloride solution Cathode-Hydrogen gas (colourless gas that extinguishes burning splint with explosion/ pop sound

AnodeOxygen gas (colourless gas that relights /rekindles glowing splint)

 

  1. Explain the difference in volume of products at the cathode and anode. Four(4) electrons donated/lost by OH ions to form 1 molecule/1volume/1mole of oxygen (O2)gas at the anode are gained/acquired/accepted by four H+(aq) ions to form 2 molecule/2volume/2mole of Hydrogen (H2)gas at the cathode.

 

The volume of Oxygen gas at the anode is half the volume of Hydrogen produced at the cathode/ The volume of Hydrogen gas at the cathode is twice the volume of Oxygen produced at the anode.

 

  1. Explain the changes in concentration of the electrolyte during electrolysis of sodium chloride solution

The concentration of dilute sodium chloride solution increases.

 

The ratio of H2 (g): O2 (g) is 2:1 as they are combined in water. Water in the electrolyte is decomposed into Hydrogen and Oxygen gases that escape as products. The concentration /moles of salt present in a given volume of sodium chloride solution continue increasing/rising.

 

 

31

 

  1. II. Dissolve about 20 g of pure sodium chloride crystals in 100cm3 of water. Place the solution in an electrolytic cell. Note the observations continuously at each electrode for 30 minutes in a fume chamber/open.

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

NaCl(aq)

 

->

->

 

OH- (aq)

 

Cl-(aq)

 

+ H+(aq)

 

+ Na+(aq)

 

 

 

  1. Name the ions in sodium chloride solution that are attracted/move to: Cathode- Na+(aq) from Sodium chloride solution (NaCl) and H+(aq) from water

(H2O)

 

Anode Cl(aq) from sodium chloride solution (NaCl) and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      2H+(aq)           +   2e     ->  H2(g)

H+ ions selectively discharged instead of Na+ ions at the cathode)

 

Anode           2Cl (aq)    ->  Cl2(g)  +  4e

(Cl ions with a higher concentration block the discharge of OH ions at the anode)

 

  1. Name the products of electrolysis of concentrated sodium chloride solution/brine

 

Cathode-Hydrogen gas (colourless gas that extinguishes burning splint with explosion/ pop sound

AnodeChlorine gas(pale green gas that bleaches damp/moist/wet litmus papers)

 

  1. Explain the difference in volume of products at the cathode and anode. Two (2) electrons donated/lost by Cl ions to form 1 molecule/1volume/1mole of Chlorine (Cl2)gas at the anode are gained/acquired/accepted by two H+(aq) ions to form 1 molecule/1volume/1mole of Hydrogen (H2)gas at the cathode.

 

The volume of Chlorine gas at the anode is equal to the volume of Hydrogen produced at the cathode/ The volume of Hydrogen gas at the cathode is equal to the volume of Chlorine produced at the anode.

 

  1. Explain the changes in concentration of the electrolyte during electrolysis of concentrated sodium chloride solution/brine

The concentration of concentrated sodium chloride solution/brine increases.

The ratio of Cl2 (g): H2 (g) is 1:1 as they are combined in water.

 

Water in the electrolyte is decomposed into only Hydrogen gas that escapes as products at cathode.

 

 

 

 

32

 

The concentration /moles of OH (aq) and Na+ ion (as NaOH) present in a given volume of electrolyte continue increasing/rising.

This makes the electrolyte strongly alkaline with high pH.

 

As the electrolysis of brine continues the concentration of Cl ions decrease and oxygen gas start being liberated at anode.

 

The electrolyte pH is thus lowered and the concentration of brine starts again increasing.

 

(ii)Electrolysis of dilute and concentrated Hydrochloric acid solution

 

  1. I. Prepare about 50cm3 of 0.05 M of dilute Hydrochloric acid in 100cm3 solution. Place the solution in an electrolytic cell. Note the observations at each electrode for 10 minutes.

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

HCl(aq)

 

->

->

 

OH- (aq)

 

Cl-(aq)

 

+ H+(aq)

 

+ H+(aq)

 

 

 

  1. Name the ions in dilute Hydrochloric acid solution that are attracted/move

to:

Cathode- H+(aq) from dilute Hydrochloric acid (HCl) and H+(aq) from water (H2O) Anode Cl(aq) from dilute Hydrochloric acid (HCl) and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      4H+(aq)           +   4e     ->  2H2(g)

H+ ions selectively discharged instead of Na+ ions at the cathode)

 

Anode           4OH (aq)   ->       H2O(l)  +O2+  4e

(4OH ions selectively discharged instead of Cl ions at the anode)

 

  1. Name the products of electrolysis of dilute Hydrochloric acid Cathode-Hydrogen gas (colourless gas that extinguishes burning splint with explosion/ pop sound

AnodeOxygen gas (colourless gas that relights /rekindles glowing splint)

 

  1. Explain the difference in volume of products at the cathode and anode. Four(4) electrons donated/lost by OH ions to form 1 molecule/1volume/1mole of oxygen (O2)gas at the anode are gained/acquired/accepted by four H+(aq) ions to form 2 molecule/2volume/2mole of Hydrogen (H2)gas at the cathode.

 

The volume of Oxygen gas at the anode is half the volume of Hydrogen produced at the cathode/ The volume of Hydrogen gas at the cathode is twice the volume of Oxygen produced at the anode.

 

 

 

 

33

 

  1. Explain the changes in concentration of the electrolyte during electrolysis of dilute Hydrochloric acid

The concentration of dilute Hydrochloric acid increases.

The ratio of H2 (g): O2 (g) is 2:1 as they are combined in water. Water in the electrolyte is decomposed into Hydrogen and Oxygen gases that escape as products. The concentration /moles of HCl present in a given volume of dilute Hydrochloric acid continue increasing/rising.

 

  1. II. Prepare about 50cm3 of 2M of Hydrochloric acid in 100cm3 solution. Place the solution in an electrolytic cell. Note the observations at each electrode for 30 minutes CautionThis experiment should be done in the open/fume chamber.

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

HCl(aq)

 

->

->

 

OH- (aq)

 

Cl-(aq)

 

+ H+(aq)

 

+ H+(aq)

 

 

 

  1. Name the ions in 2M Hydrochloric acid solution that are attracted/move to: Cathode- H+(aq) from dilute Hydrochloric acid (HCl) and H+(aq) from water (H2O) Anode Cl(aq) from dilute Hydrochloric acid (HCl) and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      4H+(aq)           +   4e     ->  2H2(g)

H+ ions selectively discharged instead of Na+ ions at the cathode)

 

Anode           2Cl (aq)   ->       Cl2+  2e

(OH ions concentration is low.Cl ions concentration is higher at the anode thus cause over voltage/block discharge of OH ions)

 

  1. Name the products of electrolysis of 2M Hydrochloric acid Cathode-Hydrogen gas (colourless gas that extinguishes burning splint with explosion/ pop sound

 

AnodeChlorine gas (Pale green gas that bleaches blue/red moist/wet/damp litmus papers)

 

  1. Explain the difference in volume of products at the cathode and anode. Two(2) electrons donated/lost by Cl ions to form 1 molecule/1volume/1mole of Chlorine (Cl2)gas at the anode are gained/acquired/accepted by two H+(aq) ions to form 1 molecule/1volume/1mole of Hydrogen (H2)gas at the cathode.

 

The volume of Chlorine gas at the anode is equal to the volume of Hydrogen produced at the cathode/ The volume of Hydrogen gas at the cathode is twice the volume of Chlorine produced at the anode.

 

34

 

  1. Explain the changes in concentration of the electrolyte during electrolysis of 2M Hydrochloric acid

The concentration of Hydrochloric acid decreases.

The ratio of H2 (g): Cl2 (g) is 1:1 as they are combined in Hydrochloric acid. Water in the electrolyte is decomposed only into Hydrogen gas that escapes as products at the cathode.

There is a net accumulation of excess OH (aq) ions in solution.

 

 

 

 

 

 

  1. Nature of electrodes used in the electrolytic cell

 

Inert electrodes (carbon-graphite and platinum) do not alter the expected products of electrolysis in an electrolytic cell. If another/different electrode is used in the electrolytic cell it alters/influences/changes the expected products of electrolysis. The examples below illustrate the influence of the nature of electrode on the products of electrolysis:

 

(i)Electrolysis of copper(II) sulphate(VI) solution I. Using carbon-graphite electrodes

 

Weigh Carbon -graphite electrodes. Record the masses of the electrodes in table I below. Place the electrodes in 1M copper(II) sulphate(VI) solution in a beaker. Set up an electrolytic cell.

 

Close the switch and pass current for about 20 minutes. Observe each electrode and any changes in electrolyte. Remove the electrodes from the electrolyte. Wash with acetone/propanone and allow them to dry. Reweigh each electrode.

 

Sample results

 

Mass of cathode before 23.4 g Mass of anode before 22.4 g
electrolysis electrolysis
Mass of cathode after 25.4 g Mass of anode after 22.4 g
electrolysis electrolysis
Brown solid deposit at the Bubbles of colourless gas
cathode after electrolysis that relight splint
Blue colour of electrolyte Blue colour of electrolyte
fades/become less blue fades /become less blue

 

 

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

 

CuSO4

 

 

(aq)

 

->

 

->

 

OH-SO4

 

(aq)

 

2-(aq)

 

+ H+(aq)

 

+  Cu2+(aq)

 

 

35

 

  1. Name the ions in 1M copper(II) sulphate(VI) solution that are attracted/move

to:

 

Cathode- Cu2+ (aq) from copper(II) sulphate(VI) solution and H+(aq) from water (H2O)

 

Anode SO42-(aq) from copper(II) sulphate(VI) solution and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      2Cu2+ (aq)    +   4e     ->  2Cu(g)

 

Cu2+ ions are lower than H+ ions in the electrochemical series therefore selectively discharged at the cathode.)

 

Anode           4OH (aq)   ->      H2O(l) + O2+  4e

(OH ions ions are higher than SO42- ions in the electrochemical series therefore selectively discharged at the cathode.))

 

  1. Name the products of electrolysis of 1M copper(II) sulphate(VI) solution Cathode-2 moles of copper metal as brown solid coat

 

AnodeOxygen gas (Colourless gas that relights /rekindles glowing splint)

 

  1. Explain the changes that take place at the cathode and anode.

 

Four(4) electrons donated/lost by OH ions to form 1 molecule/1volume/1mole of Oxygen (O2)gas at the anode are gained/acquired/accepted by two Cu2+(aq) ions to form 2 moles of brown copper solid that deposit itself at the cathode.

 

The moles of oxygen gas at the anode is equal to the moles of copper produced at the cathode

 

  1. Explain the changes in electrolyte during electrolysis of 1M copper (II) sulphate(VI) solution.

 

(i)The pH of copper(II) sulphate(VI) solution lowers/decreases. The salt becomes more acidic. Water in the electrolyte is decomposed only into Oxygen gas (from the OH ions) that escapes as products at the anode. There is a net accumulation of excess H+ (aq) ions in solution. This makes the electrolyte strongly acidic with low pH.

 

  • Cu2+ (aq) ions are responsible for the blue colour of the electrolyte/ copper(II)

sulphate (VI) solution. As electrolysis continues, blue Cu2+ (aq) ions gain electrons to form brown Copper. The blue colour of electrolyte therefore fades/become less

blue.

 

(iii)Copper is deposited at the cathode. This increases the mass of the cathode.OHions that produce Oxygen gas at anode come from water. Oxygen escapes out/away without increasing the mass of anode.

 

36

 

  1. Using copper electrodes

 

Weigh clean copper plates electrodes. Record the masses of the electrodes in table I below. Place the electrodes in 1M copper(II) sulphate(VI) solution in a beaker. Set up an electrolytic cell.

 

Close the switch and pass current for about 20 minutes. Observe each electrode and any changes in electrolyte. Remove the electrodes from the electrolyte. Wash with acetone/propanone and allow them to dry. Reweigh each electrode.

 

Sample results

 

Mass of cathode before 23.4 g Mass of anode before 22.4 g
electrolysis electrolysis
Mass of cathode after 25.4 g Mass of anode after 20.4 g
electrolysis electrolysis
Brown solid deposit at the Anode decrease
cathode after electrolysis insize/erodes/wear off
Blue colour of electrolyte Blue colour of electrolyte
remain blue remain blue

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

 

CuSO4

 

 

(aq)

 

->

 

->

 

OH-SO4

 

(aq)

 

2-(aq)

 

+ H+(aq)

 

+  Cu2+(aq)

 

 

  1. Name the ions in 1M copper(II) sulphate(VI) solution that are attracted/move

 

to:

 

Cathode- Cu2+ (aq) from copper(II) sulphate(VI) solution and H+(aq) from water (H2O)

 

Anode SO42-(aq) from copper(II) sulphate(VI) solution and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      Cu2+ (aq)       +   2e   ->  Cu(s)

Cu2+ ions are lower than H+ ions in the electrochemical series therefore selectively discharged at the cathode.)

 

Anode Cu (s)  ->   Cu2+(aq)+  2e
(Both OH ions  and  SO4 2- ions  move to the anode but none is discharged. The

 

copper anode itself ionizes/dissolves/dissociate because less energy is used to remove an electron/ionize /dissociate copper atoms than OH ions.

 

 

 

37

 

  1. Name the products of electrolysis of 1M copper(II) sulphate(VI) solution using copper electrodes.

 

Cathode-1 moles  of copper metal as brown solid coat (Cathode increase/deposits)

 

Anode-Anode erodes/decrease in size

 

  1. Explain the changes that take place during the electrolytic process (i)Cathode

 

-Cu2+ ions are lower than H+ ions in the electrochemical series therefore selectively discharged at the cathode. Cu2+ ions have greater tendency to accept/gain/acquire electrons to form brown copper atoms/solid that deposit itself and increase the mass/size of the cathode.The copper deposited at the cathode is pure

 

-H+ ions accumulate around the cathode. Electrolyte thus becomes strongly acidic around the cathode.

 

-Cu2+ ions in solution are responsible for the blue colour of electrolyte. Blue colour of electrolyte fade around the cathode.

 

(ii)Anode

 

Copper atom at the anode easily ionizes to release electrons. The anode therefore keeps decreasing in mass/eroding. The amount of copper that dissolve/erode is equal to the mass of copper deposited. This is called electrode ionization.

 

Electrode ionization is where the anode erodes/decrease and the cathode deposits/increase during electrolysis. The overall concentration of the electrolyte remains constant

 

14.In industries electrolysis has the following uses/applications:

 

(a)Extraction of reactive metals from their ores.

 

Potassium, sodium ,magnesium, and aluminium are extracted from their ores using electrolytic methods.

 

(b)Purifying copper after exraction from copper pyrites ores.

 

Copper obtained from copper pyrites ores is not pure. After extraction, the copper is refined by electrolysing copper(II)sulphate(VI) solution using the impure copper as anode and a thin strip of pure copper as cathode. Electrode ionization take place there:

 

 

(i)At the cathode; (ii)At the anode;

 

Cu2+ (aq) + 2e -> Cu(s) (Pure copper deposits on the strip Cu(s) ->Cu2+ (aq) + 2e (impure copper erodes/dissolves)

 

 

 

(c)Electroplating

 

The label EPNS(Electro Plated Nickel Silver) on some steel/metallic utensils mean they are plated/coated with silver and/or Nickel to improve their appearance(add their aesthetic value)and prevent/slow corrosion(rusting of iron). Electroplating is the process of coating a metal with another metal using an electric current. During electroplating,the cathode is made of the metal to be coated/impure.

 

38

 

Example:

During the electroplating of a spoon with silver

 

(i)the spoon/impure is placed as the cathode(negative terminal of battery) (ii)the pure silver is placed as the anode(positive terminal of battery) (iii)the pure silver erodes/ionizes/dissociates to release electrons:

Ag(s) ->Ag+ (aq)     +  e             (impure silver erodes/dissolves)

 

  • silver (Ag+)ions from electrolyte gain electrons to form pure silver deposits / coat /cover the spoon/impure

Ag+ (aq)       +   e  ->Ag(s)    (pure silver deposits /coat/cover on spoon)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

15.The quantitative amount of products of electrolysis can be determined by applying Faradays 1st law of electrolysis.

 

Faradays 1st law of electrolysis states that the mass/amount of substance liberated/produced/used during electrolysis is directly proportional to the quantity of of electricity passed/used.

 

(a)The SI unit of quantity of electricity is the coulomb(C). The coulomb may be defined as the quantity of electricity passed/used when a current of one ampere flow for one second.i.e;

 

1Coulomb = 1 Ampere x 1Second The Ampere is the SI unit of current(I)

The Second is the SI unit of time(t) therefore;

 

 

39

 

Quantity of electricity(in Coulombs) = Current(I)  x  time(t)

Practice examples

 

  1. A current of 2 amperes was passed through an electrolytic cell for 20 minutes. Calculate the quantity of electric charge produced.

 

Working:

Quantity of electricity(in Coulombs) = Current(I)  x  time(t)
Substituting /converting time to second = 2  x (20 x 60)
= 2400 C

 

  1. A current of 2 amperes was passed through an electrolytic.96500 coulombs of charge were produced. Calculate the time taken.

 

Working:

Time(t) in seconds = Quantity of electricity(in Coulombs)
Current(I) in amperes
Substituting = 96500
2
  • 48250 seconds

 

 

  1. 96500 coulombs of charge were produced after 10 minutes in an electrolytic cell . Calculate the amount of current used.

Working:

Current(I) in amperes = Quantity of electricity(in Coulombs)
Time(t) in seconds
Substituting/converting time to second= 96500
10 x 60
= 160.8333 Amperes

 

(b)The quantity of electricity required for one mole of electrons at the anode/cathode is called the Faraday constant(F). It is about 96500 Coulombs.i.e

 

The number of Faradays used /required is equal to the number of electrons used at cathode/anode during the electrolytic process. e.g.

 

Cu2+ require to gain 2 moles of electrons=2 Faradays =2 x 96500 coulombs of electricity at the cathode.

 

Al3+ require to gain 3 moles of electrons=3 Faradays =3 x 96500 coulombs of electricity at the cathode

 

Na+ require to gain 1 moles of electrons=1 Faradays =1 x 96500 coulombs of electricity at the cathode

 

2H+ require to gain 2 moles of electrons=2 Faradays =2 x 96500 coulombs of electricity at the cathode to form 1molecule of hydrogen gas

 

2O2 require to lose/donate 4 moles of electrons=4 Faradays =4 x 96500 coulombs of electricity at the anode to form 1molecule of Oxygen O2 gas.

 

4OH require to lose/donate 4 moles of electrons=4 Faradays =4 x 96500 coulombs of electricity at the anode to form 1molecule of Oxygen gas and 2 molecules of water.

 

40

 

(c)The mass/amount of products at the cathode/anode is related to the molar mass of the substance and/or the volume of gases at standard/room temperature and pressure as in the below examples:

 

Practice examples

 

1.Calculate the mass of copper deposited at the cathode when a steady current of

 

4.0 amperes is passed through copper(II)sulphate(VI) for 30 minutes in an electrolytic cell. (Cu=63.5, 1F = 96500C)

 

Working:
Quantity of electricity(in Coulombs) = Current(I)  x  time(t)
Substituting /converting time to second = 4  x (30 x 60)
= 7200 C
Equation at the cathode:  Cu2+ (aq) +  2e
->  Cu(s)

 

2 mole of electrons = 2 Faradays = 2 x 96500 C produce a mass =molar mass of copper thus;

2  x  96500C -> 63.5 g
72000C -> 7200  x  63.5 2.3689 g of copper
2  x 96500

 

2.a)If 3.2 g of Lead were deposited when a current of 2.5 amperes was passed through an electrolytic cell of molten Lead(II)bromide for 20 minutes, determine the Faraday constant.(Pb = 207)

Working:
Quantity of electricity (in Coulombs) = Current(I)  x  time(t)
Substituting /converting time to second = 2.5 x (20 x 60)
= 3000 C
If 3.2g of Lead -> 3000C
Then 207 g of Lead -> 207 x 3000 = 194062.5 C
3.2
Equation at the cathode:  Pb2+ (l) +  2e -> Pb(l)
From the equation: 2 moles of electrons = 2 Faradays =  194062.5 C
1mole of electrons =  1 Faraday  => 194062.5  =  97031.25 C
2

 

b)What is the volume of bromine vapour produced at the anode at room temperature(1mole of gas at room temperature and pressure = 24000cm3)

 

Method 1
Equation at the anode: Br (l) -> Br2(g)  +  2e
41

 

From the equation: 2 moles of electrons = 2 Faradays = 194062.5 C -> 24000cm3 3000 C -> 3000 x 24000

 

194062.5

=371.0145cm3
Method 2
Br (l) -> Br2(g)  +  2e
Equation at the anode:
Mole ratio of products at Cathode: anode = 1:1
Moles of Lead at cathode = 3.2  = 0.0155moles = moles of Bromine
207
1 moles of bromine vapour  -> 24000cm3
0.0155moles  of Bromine -> 0.0155 x 24000  = 372 cm3
Method 3
Br (l) -> Br2(g)  +  2e
Equation at the anode:
Ratio of Faradays used to form products at Cathode: anode = 2:2
=>  2 x 97031.25 C produce 24000cm3 of bromine vapour
Then: 3000 C  ->  3000 x 24000cm3 = 371.0145cm3
2 x 97031.25

 

3.What mass of copper remain from 2.0 at the anode if a solution of copper(II)sulphate(VI) is electrolysed using a current of 1 ampere flowing through an electrolytic cell for 20 minutes.(Cu= 63.5, 1Faraday = 96487 coulombs)

 

Working:

 

Quantity of electricity (in Coulombs) = Current(I)  x  time(t)
Substituting /converting time to second = 1  x (20 x 60)
= 1200 C
Equation at the cathode:Cu2+ (aq)  +
2e  -> Cu(s)

 

2 mole of electrons = 2 Faradays = 2 x 96500 C erode/dissolve a mass =molar mass of copper thus;

2  x  96500C -> 63.5 g
1200C -> 1200  x  63.5  =  0.3948g of copper deposited
2  x 96500

Mass of copper remaining = Original mass mass dissolved/eroded

 

=>  2.0 -0.3948      =         1.6052 g of copper remain

 

  1. Calculate the current passed if a mass of 0.234 g of copper is deposited in 4 minutes during electrolysis of a solution of copper (II)sulphate(VI).

 

(Cu= 63.5  ,1F = 96500C)

Working:

 

Equation at the cathode:        Cu(s)   ->   Cu2+ (aq)   +  2e

 

2 mole of electrons = 2 Faradays = 2 x 96500 C produce a mass =molar mass of copper thus;

 

42

 

63.5 g    ->   2  x 96500C
0.234 g  ->  0.234 x 2 x 96500 = 711.2126 C
63.5
Current(I) in amperes = Quantity of electricity(in Coulombs)
Time(t) in seconds
Substituting/converting time to second= 711.2126 C
4x 60
  • 9634 Amperes

 

  1. (a)What quantity of electricity will deposit a mass of 2.43 g of Zinc during electrolysis of a solution of Zinc (II)sulphate(VI).

 

(Zn= 65  ,1F = 96500C)

Working:

Equation at the cathode:        Zn2+ (aq)   +  2e  -> Zn(s)

 

2 mole of electrons = 2 Faradays = 2 x 96500 C erode/dissolve a mass =molar mass of Zinc thus;

 

65 g -> 2  x  96500
2.43 g -> 2.43 x 2  x  96500   =  7215.2308 C
65

 

(b)Calculate the time (in minutes) it would take during electrolysis of the solution of Zinc (II)sulphate(VI) above if a current of 4.0 Amperes is used.

 

Time(t) in seconds = Quantity of electricity(in Coulombs)
Current(I) in amperes
Substituting =  7215.2308 = 1803.8077 seconds  =  30.0635 minutes
4 60

 

6.When a current of 1.5 amperes was passed through a cell containing M3+ ions of metal M for 15 minutes, the mass at cathode increased by 0.26 g.(Faraday constant = 96500C

 

  1. a) Calculate the quantity of electricity used.

Quantity of electricity (in Coulombs)  = Current(I)  x  time(t)

Substituting /converting time to second             = 1.5  x (15 x 60)

= 1350 C

 

  1. Determine the relative atomic mass of metal M

Equation at the cathode:M3+ (aq)  +  3e  -> M(s)

1350 C of electricity    ->  0.26 g  of metal M

 

3 mole of electrons = 3 Faradays = 3 x 96500 C produce a mass =molar mass of  M

thus;

RAM of M    = 0.26 g  x  3  x    96500    =    55.7556(No units)

1350

 

7.An element P has a relative atomic mass 88.When a current of 0.5 amperes was passed through fused chloride of P for 32 minutes and 10seconds ,0.44 g

 

43

 

of P was deposited at the cathode. Determine the charge on an ion of P(Faraday constant = 96500C)

Working:
Quantity of electricity (in Coulombs) = Current(I)  x  time(t)
Substituting /converting time to second = 0.5  x ((32 x 60) + 10)
= 965C
0.44 g of metal Pare deposited by 965C
88g of of metal Pare deposited by:  88  x 965= 193000 C
0.44
96500 C  = 1 mole of electrons = 1 Faradays = single charge
193000 C ->   193000 = 2  moles/Faradays/charges => symbol of ion = P2+
96500

 

  1. During purification of copper by electrolysis 1.48 g of copper was deposited when a current was passed through aqueous copper (II)sulphate(VI) for 2 ½
hours. Calculate the amount of current that was passed. (Cu= 63.5
,1F = 96500C)
Working:
Equation at the cathode: Cu2+ (aq)  +  2e-> Cu(s)

2 mole of electrons = 2 Faradays = 2 x 96500 C produce a mass =molar mass of

copper  thus;
63.5 g    ->   2  x 96500C
1.48 g  ->  1.48 x 2 x 96500 = 4255.1181 C
63.5
Current(I) in amperes = Quantity of electricity(in Coulombs)
Time(t) in seconds
Substituting/converting time to second= 4255.1181C
(( 2  x 60) + 30) x60
  • 4728 Amperes

 

  1. Practically Faraday 1st law of electrolysis can be verified as

 

Verifying Faraday 1st law of electrolysis

Procedure.

 

Weigh clean copper plates electrodes. Record the masses of the electrodes in table I below. Place the electrodes in 1M copper(II) sulphate(VI) solution in a beaker. Set up an electrolytic cell.

 

Close the switch and pass a steady current of 2 amperes by adjusting the rheostat for exactly 20 minutes.Remove the electrodes from the electrolyte. Wash with acetone/ propanone and allow them to dry. Reweigh each electrode.

 

Sample results

 

Mass of cathode before 7.00 g Mass of anode before 7.75 g
electrolysis electrolysis
Mass of cathode after 8.25 g Mass of anode after 6.50 g
44

 

 

electrolysis electrolysis
Change in mass at 1.25 g Change in mass at anode 1.25 g
cathode after electrolysis after electrolysis

 

Answer the following questions:

 

  1. Write the equation for the decomposition of the electrolytes during the electrolytic process.

 

 

H2O(l)

CuSO4

 

 

(aq)

 

->

->

 

OH-SO4

 

(aq)

 

2-(aq)

 

+ H+(aq)

+  Cu2+(aq)

 

 

  1. Name the ions in 1M copper(II) sulphate(VI) solution that are attracted/move

to:

 

Cathode- Cu2+ (aq) from copper(II) sulphate(VI) solution and H+(aq) from water (H2O)

Anode SO42-(aq) from copper(II) sulphate(VI) solution and OH (aq) from water (H2O)

 

III. Write the equation for the reaction during the electrolytic process at the:

Cathode      Cu2+ (aq)       +   2e   ->  Cu(s)

 

Cu2+ ions are lower than H+ ions in the electrochemical series therefore selectively discharged at the cathode.)

 

Anode           Cu (s)   ->    Cu2+(aq)      +  2e

(Both OH ions and SO42- ions move to the anode but none is discharged. The copper anode itself ionizes/dissolves/dissociate as less energy is used to remove an electron/ionize /dissociate copper atoms than OH ions.

 

  1. Name the products of electrolysis of 1M copper(II) sulphate(VI) solution

using copper electrodes.

Cathode-1.25 g  of copper metal as brown solid coat/deposits

 

Anode-1.25 g  of copper metal erodes/decrease in size

 

  1. (i)How many moles of electrons are used to deposit/erode one mole of copper metal at the cathode/anode?

From the equation at anode/cathode= 2 moles

 

(ii)How many Faradays are used to deposit/erode one mole of copper metal at the cathode/anode?

From the equation at anode/cathode : 2 moles  = 2 Faradays

 

(iii)Calculate the quantity of electric charge used

 

Working:

 

45

 

 

Quantity of electricity (in Coulombs) Substituting /converting time to second

 

  • Current(I) x  time(t)
  • 2 x  20 x 60
    • 2400C

 

 

 

  1. (i) Calculate the quantity of electricity required to deposit/erode one mole of copper at the cathode/anode(Cu=63.5)
Since 1.25 g of copper -> 2400C
Then 63.5 g (1mole of copper) -> 63.5  x  2400 =  121920 C
1.25

 

(ii)Determine the Faraday constant from the results in V(i) above From the equation at;

Cathode Cu2+ (aq) +  2e  ->  Cu(s)
Anode Cu (s)  -> Cu2+(aq)+  2e
2 moles = 2 Faradays -> 121920 C
1 moles = 1 Faradays -> 121920  = 60960 C
2

 

  • The faraday constant obtained above is far lower than theoretical.Explain

 

-high resistance of the wires used.

-temperatures at 25oC were not kept constant

-plates/electrodes used were not made of pure copper

-plates/electrodes used were not thoroughly clean copper

 

Further practice

 

1.An element P has a relative atomic mass of 88. When a current of 0.5 amperes was passed through the fused chloride of P for 32 minutes and 10 seconds, 0.44g of P were deposited at the cathode. Determine the charge on an ion of P. (1 faraday = 96500 Coulombs).

 

2.During electrolysis of aqueous copper (II) sulphate, 144750 coulombs of electricity were used. Calculate the mass of copper metal that was obtained

(Cu = 64       ;1 Faraday = 96500 coulombs)         ( 3 mks)

 

3.A nitrate of a metal M was electrolysed .1.18 g of metal was deposited when a current of 4 ampheres flow for 16 minutes.Determine the formula of the sulphate(VI)salt of the metal.

 

(Faraday constant = 96500 , RAM of X = 59.0)

Working

 

Q  = It     =>( 4 x 16 x 60) = 3840 C

1.18 g of X => 3840 C

59.0 g =>       59.0 x 3840  = 192000 C

1.18

96500 C = 1Faraday

192000 C=  192000 C x1   =      2F thus charge of M = M2+

96500 C

Valency of  M  is 2 thus formula of sulphate(VI)salt  MSO4

 

46

 

  1. Below is the results obtained when a current of 2.0ampheres is passed through copper(II)sulphate(VI)solution for 15 minutes during electrolysis using copper electrode.
Initial mass of cathode = 1.0 g
Final mass of cathode = 1.6 g

Change in mass of cathode  =  0.60 g

(i)Determine  the change in mass at the anode. Explain your answer.

Mass decrease  = 0.6g.

Electrode ionization take place where the cathode increase in mass form the erosion of the anode

 

(ii)Calculate the quantity of electricity required to deposit one mole of copper.(Cu =63.5) Q =It => 2 x 15 x 60 = 1800 coulombs

 

Method 1
0.60 g  of copper ->1800 coulombs
63.5 g -> 63.5 x 1800 190500 Coulombs
0.60
Method 2
Moles of Copper = Mass => 0.60  =  9.4488 x10 -3 moles
Molar mass 63.5
9.4488 x10 -3 moles -> 1800 coulombs
1 Mole -> 1 x 1800 coulombs    =  190500.381 coulombs
9.4488 x10 -3 moles

 

(iii)Determine the oxidation number of copper produced at the cathode and hence the formula of its nitrate (V)salt (1 Faraday = 96500 Coulombs)

 

96500 Coulombs -> 1 Faraday
190500.381 coulombs -> 190500.381 coulombs  x 1
96500 Coulombs
= 1.9741 Faradays   => 2F(whole number)
Charge of copper = 2+ Oxidation number
= 2 hence chemical formula of nitrate (V)salt = Cu (NO 3)2
=>  Valency of copper

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

47

 

22.0.0 METALS (20 LESSONS)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

a)Introduction to metals

 

The rationale of studying metals cannot be emphasized.Since ages, the world over, metals like gold and silver have been used for commercial purposes.

 

The periodicity of alkali and alkaline earth metals was discussed in year 2 of secondary school education. This topic generally deals with:

 

(a)Natural occurrence of the chief ores of the most useful metals for industrial /commercial purposes.

 

(b)Extraction of these metals from their ores for industrial/ commercial purposes.

(c)industrial/ commercial uses of these metals.

(d)main physical and chemical properties /characteristic of the metals.

 

The metals given detailed emphasis here are; Sodium, Aluminium, Iron, Zinc, Lead and Copper.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

1

 

The main criteria used in extraction of metals is based on its position in the electrochemical/reactivity series and its occurrence on the earths crust.

 

Position on the earth’s crust

 

 

 

 

 

 

 

 

If near the surface ,open

 

cast mining / quarrying is

 

used

 

 

 

 

 

 

 

If deep on the earth’s crust

 

deep mining is used

 

 

 

 

If the ore is low grade  oil, water, and air

 

is  blown  forming a froth(froth flotation)

 

to concentrate

 

Electrolysis of  the ore is used for

 

reactive metals; Potassium, Sodium,

 

Magnesium, Calcium, Aluminium

 

The ore first roasted if it is a carbonate or

 

sulphide  of Zinc, Iron, Tin, Lead, and Copper

 

to form the oxide

 

 

 

 

The oxide is reduced using carbon/ carbon(II)

 

oxide in a furnace if it is made of Zinc ,Tin,

 

Lead ,Copper and Iron

 

1.SODIUM

 

  1. Natural occurrence Sodium naturally occurs as:

 

(i)Brine-a concentrated solution of sodium chloride(NaCl(aq)) in salty seas and oceans.

(ii)Rock salt-solid sodium chloride(NaCl(s)

(iii)Trona-sodium sesquicarbonate(NaHCO3.Na2CO3.2H2O) especially

 

in lake Magadi in Kenya.

 

(iv)Chile saltpeter-sodium nitrate(NaNO3)

 

 

2

 

b)(i)

 

Extraction of Sodium from brine/Manufacture of Sodium hydroxide/The flowing mercury cathode cell/ TheCaster-Keller process

 

I.Raw materials

 

  • Brine-concentrated solution of sodium chloride (NaCl (aq)) from salty seas and oceans.

(ii)Mercury

 

  • Water from river/lakes

 

  1. Chemical processes

 

Salty lakes, seas and oceans contain large amount of dissolved sodium chloride (NaCl (aq)) solution.

 

This solution is concentrated to form brine which is fed into an electrolytic chamber made of suspended Carbon graphite/titanium as the anode and a continuous flow of Mercury as the cathode.Note

 

Mercury is the only naturally occurring known liquid metal at room temperature and pressure

 

Questions

 

  1. Write the equation for the decomposition of the electrolyte during the electrolytic process.

 

 

H2O(l)

NaCl(aq)

 

 H+(aq)

+

 

+

 

+

 

OH(aq)

 

 

 

  1. Name the ions present in brine that moves to the: (i)Mercury cathode; H+(aq) , Na+(aq) (ii)Titanium/graphite; OH(aq), Cl(aq)

 

  • Write the equation for the reaction that take place during the electrolytic process at the;

Cathode;     2Na+(aq)  +           2e                                 2Na(s)

Anode;         2Cl (aq)                                 Cl2(g)      +  2e

Note

(i)Concentration of 2Cl(aq) ions is higher than OH ions causing overvoltage thus blocking OH ions from being discharged at the anode.

 

(ii)Concentration of Na+(aq) ions is higher than H+ ions causing overvoltage thus blocking H+ ions from being discharged at the cathode.

 

3

 

  1. Name the products of electrolysis in the flowing mercury-cathode cell.

 

(i)Mercury cathode; Sodium metal as grey soft metal/solid

 

 

(ii)Titanium/graphite; Chlorine gas as a pale green gas that turns moist blue/red litmus papers red then bleaches both. Chlorine gas is a very useful by-product in;

 

(i)making (PVC)polyvinylchloride(polychloroethene) pipes. (ii)chlorination/sterilization of water to kill germs. (iii)bleaching agent

 

(iv)manufacture of hydrochloric acid.

 

Sodium produced at the cathode immediately reacts with the mercury at the cathode forming sodium amalgam(NaHg) liquid that flow out of the chamber.

 

Na(s)  +   Hg(l)    Na Hg (l)

 

Sodium amalgam is added distilled water and reacts to form sodium hydroxide solution, free mercury and Hydrogen gas.

2Na Hg (l)  + 2H2O(l)  2NaOH (aq) +  2Hg(l)  +  H2(g)

 

Hydrogen gas is a very useful by-product in;

 

(i)making ammonia gas in the Haber process (ii)manufacture of hydrochloric acid (iii)in weather balloons to forecast weather (iv)as rocket fuel

 

As the electrolysis of brine continues, the concentration of Cl-ions decreases and oxygen gas start being liberated. Continuous feeding of the electrolyte is therefore very necessary.

 

III.Uses of sodium hydroxide

The sodium hydroxide produced is very pure and is used mainly in:

 

(i)Making soapy and soapless detergents.

(ii)making cellulose acetate/rayon

 

 

  1. Diagram showing the Manufacture of Sodium hydroxide from the flowing Mercury-cathode cell.

 

 

 

 

 

 

4

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

3

 

 

 

  1. Environmental effects of Manufacture of Sodium hydroxide from the flowing Mercury-cathode cell.

 

1.Most of the Mercury used at the cathode is recycled ;

(i)to reduce the cost because mercury is expensive

 

(ii)to reduce pollution because mercury kills marine life.

(iii)because it causes chromosomal/genetic mutation to human beings.

 

2.Chlorine produced at the anode;

(i)has a pungent irritating smell that causes headache to human beings.

 

(ii)bleaches any wet substance.

 

(iii)dissolves water to form both hydrochloric acid and chloric(I)acid Both cause marine pollution and stomach upsets.

 

b)(ii)

Extraction of sodium from rock salt/The Downs cell/process

 

5

 

  1. Raw materials

(i)Rock salt/solid sodium chloride

 

(ii)calcium(II)chloride

 

  1. Chemical processes.

 

Rock salt/ solid sodium chloride is heated to molten state in a chamber lined with fire bricks on the outside.

 

Sodium chloride has a melting point of about 800oC. A little calcium (II) chloride is added to lower the melting point of the electrolyte to about 600oC. The molten electrolyte is the electrolyzed in a carbon graphite anode suspended at the centre and surrounded by steel cathode.

 

Questions

 

  1. Write the equation for the decomposition of the electrolyte during the electrolytic process.

NaCl(l)    Na+(l)            +           Cl(l)

Note: In absence of water, the ions are in liquid state.

 

  1. Name the ions present in molten rock salt that move to the; (i)Steel cathode -Na+(l)

(ii)Carbon graphite anode- Cl(l)

 

  • Write the equation for the reaction that take place during the electrolytic process at the;

 

(i)Steel cathode

2Na+(l)        +             2e  2Na(l)

(ii)Carbon graphite anode

2Cl(l)   Cl2(g)        +        2e

 

  1. Name the products of electrolysis in the Downs cell at; (i)Cathode:

 

Grey solid Sodium metal is less dense than the molten electrolyte and therefore float on top of the cathode to be periodically tapped off.

 

(ii)Anode:

 

Pale green chlorine gas that turns moist/damp/wet blue/red litmus papers red then bleaches/decolorizes both. Chlorine gas is again a very useful by-product in;

 

(i)making (PVC)polyvinylchloride(polychloroethene) pipes. (ii)chlorination/sterilization of water to kill germs. (iii)bleaching agent

 

(iv)manufacture of hydrochloric acid.

 

6

 

  • steel diaphragm/gauze is suspended between the electrodes to prevent recombination of sodium at the cathode and chlorine gas at the anode back to sodium chloride.

 

  • Diagram showing the Downs cell/process for extraction of sodium

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

  1. Uses of sodium.

 

1.Sodium vapour is used as sodium lamps to give a yellow light in street lighting.

 

2.Sodium is used in making very useful sodium compounds like; (i)Sodium hydroxide(NaOH)

 

(ii)Sodium cyanide(NaCN)

(iii)Sodium peroxide(Na2O2)

(iv)Sodamide(NaNH2)

3.An alloy of Potassium and Sodium is used as coolant in nuclear reactors.

 

  1. Environmental effects of Downs cell.

 

1.Chlorine produced at the anode;

 

7

 

(i)has a pungent irritating smell that causes headache to human beings.

(ii)bleaches any wet substance.

 

(iii)dissolves water to form both hydrochloric acid and chloric(I)acid Both cause marine pollution and stomach upsets.

 

2.Sodium metal rapidly react with traces of water to form alkaline Sodium hydroxide(NaOH(aq))solution. This raises the pH of rivers/lakes killing aquatic lifein case of leakages.

 

  1. Test for presence of Na.

 

If a compound has Na+ ions in solid/molten/aqueous state then it changes a non-luminous clear/colourless flame to a yellow coloration but does not burn

 

Experiment

 

 

Scoop a portion of sodium chloride crystals/solution in a clean metallic spatula. Introduce it to a clear /colourless Bunsen flame.

 

 

Observation Inference
Yellow coloration Na+

 

Practice

 

(i)Calculate the time taken in hours for 230kg of sodium to be produced in the Downs cell when a current of 120kA is used.

 

(ii)Determine the volume of chlorine released to the atmosphere. (Na=23.0),Faraday constant=96500C.I mole of a gas =24dm3 at r.t.p) Working:

 

Equation at the cathode:

2Na+ (l)   +  2e  -> 2Na(l)

 

2 mole of electrons = 2 Faradays = 2 x 96500 C deposits a mass = molar mass of Na = 23.0g thus;

 

23.0 g -> 2  x  96500 C
(230 x 1000)g ->230 x 1000 x 2  x  96500
23
= 1,930,000,000 / 1.93 x 10 9C

 

Time(t) in seconds       =  Quantity of electricity

 

Current(I) in amperes

 

8

 

Substituting

  • 1,930,000,000 / 1.93 x 10 9C

 

120 x 1000A

  • 16,083,3333seconds / 268.0556 minutes

 

  • 4676hours

 

Volume of Chlorine

Method 1

 

Equation at the anode:

2 Cl  (l) -> Cl2(g)  +  2e

 

2 moles of electrons = 2 Faradays  =2 x 96500C

2 x 96500C  -> 24dm3

9                                                                              9

1,930,000,000 / 1.93 x 10  C->1,930,000,000 / 1.93×10  C x 24

 

Volume of Chlorine = 240,000dm3 /2.4 x 105dm3

 

Method 2
Equation at the anode: Cl (l) -> Cl2(g)  +  2e
Mole ratio of products at Cathode: anode = 1:1

 

Moles of  sodium  at cathode =(230 x 1000 )g= 10,000moles

23
10,000moles of  Na   = 10,000moles moles of Chlorine
1 moles of Chlorine gas =  24000cm3
10,000moles of Chlorine- >  10000 x 24

=240,000dm3 / 2.4x 105dm3

 

Method 3

Equation at the anode: Cl (l) -> Cl2(g)  +  2e

Ratio of Faradays of products at Cathode: anode = 2:2

 

=> 2 x 96500C produce 24000cm3 of chlorine gas Then: 1,930,000,000 / 1.93 x 10 9C ->

9

1,930,000,000 / 1.93 x 10  C x24  = 240,000dm3

 

 

 

 

(iij)The sodium metal produced was reacted with water to form 25000dm3 solution in a Caster-Keller tank.

 

9

 

(a)Calculate the concentration of the resulting solution in moles per litre. (b)The volume of gaseous products formed at s.t.p(1 mole of gas =22.4 dm3 at s.t.p)

 

Chemical equation at Caster-Keller tank

2Na(s)  +  2H2O(l) -> 2NaOH(aq)  +  H2 (g)

 

Mole ratio Na:NaOH =  2 : 2 => 1:1

 

Moles Na =10000moles=10000moles of NaOH 25000dm3 ->10000moles of NaOH

1dm3 -> 10000  x 1   =  0.4M /  0.4 moles/dm3
25000

 

Mole ratio Na: H2 (g) =  2 : 1

Moles Na = 10000moles = 5000moles of H2 (g) Volume of H2 (g) = moles x molar gas volume at s.t.p

=> 5000moles x 22.4 dm3

 

=120,000dm3

 

 

(iv)The solution formed was further diluted with water for a titration experiment. 25.0 cm3 of the diluted solution required 20.0cm3 of 0.2M sulphuric(VI)acid for complete neutralization. Calculate the volume of water added to the diluted solution before titration.

 

Chemical equation

2NaOH(aq)  +  H2SO4(aq)  -> Na2SO4(aq)  +  H2O(l)

Moles ratio NaOH : H2SO4 =  2:1
Moles ratio  H2SO4 = molarity x volume => 0.2M  x  20
1000 1000
=4.0 x 10-3 moles

 

Moles NaOH = 2 x 4.0 x 10-3 moles= 8.0 x 10-3 moles Molarity of NaOH= Moles x 1000=> 8.0 x 10-3 moles x 1000

volume                                              25

=0.16 molesdm-3 /M

 

Volume used during dilution

C1V1 = C2V2   =>         0.4M   x  V1  =  0.16 M  x  25

 

10

 

= 0.16 M x 25     = 10cm3

0.4

 

 

  • Below is a simplified diagram of the Downs Cell used for the manufacture of sodium. Study it and answer the questions that follow

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(i)What material is the anode made of? Give a reason (2 mks) Carbon graphite/Titanium

 

This because they are cheap and inert/do not influence/affect the products of electrolysis

 

  • What precaution is taken to prevent chlorine and sodium from re-combination? ( 1 mks)

 

Using a steel gauze/diaphragm separating the cathode from anode

 

  • Write an ionic equation for the reaction in which chlorine gas is

 

formed( 1mk)

 

2Cl(l)          ->        Cl2(g)       +            2e

 

  • In the Downs process, (used for manufacture of sodium), a certain

 

salt is added to lower the melting point of sodium chloride from about 8000C to about 6000C.

(i) Name the salt that is added                                (1mk)

Calcium chloride

 

  • State why it is necessary to lower the temperature(1mk) To reduce the cost of production

 

11

 

  • Explain why aqueous sodium chloride is not suitable as an electrolyte for the manufacture of sodium in the Downs process( 2mk)

 

The sodium produced react explosively/vigorously with water in the aqueous sodium chloride

 

  • Sodium metal reacts with air to form two oxide. Give the formulae of

two oxides ( 1mk)

Na2O Sodium oxide(in limited air) Na2O2 Sodium peroxide(in excess air)

 

2.ALUMINIUM

 

a)Natural occurrence

 

Aluminium is the most common naturally occurring metal. It makes 7% of the earths crust as:

(i)Bauxite ore- Hydrated aluminium oxide(Al2O3.2H2O)

(ii)Mica ore-Potassium aluminium silicate(K2Al2Si6O16)

(iii)China clay ore- aluminium silicate (Al2Si6O16)

(iv)Corrundum-Anhydrous aluminium oxide(Al2O3)

 

b)Extraction of aluminium from Bauxite/Halls cell/process)

 

The main ore from which aluminium is extracted is Bauxite ore- hydrated aluminium oxide(Al2O3.2H2O).

 

The ore is mined by open-caste mining method/quarrying where it is scooped together with silica/sand/silicon(IV)oxide (SiO2) and soil/ iron(III)oxide

(Fe2O3) as impurities.

 

The mixture is first dissolved in hot concentrated sodium/potassium hydroxide solution.

 

The alkalis dissolve both bauxite and silicon(IV)oxide.

 

This is because bauxite is amphotellic while silicon(IV)oxide is acidic. Iron(III)oxide (Fe2O3) is filtered of /removed as a residue. Carbon(IV)oxide is bubbled into the filtrate to precipitate aluminium (III) hydroxide (Al(OH)3) as residue.

 

The aluminium (III) hydroxide (Al(OH)3) residue is filtered off. Silicon (IV)oxide remain in the solution as filtrate. Aluminium (III) hydroxide (Al(OH)3) residue is then heated to form pure aluminium (III)oxide(Al2O3)

 

 

2Al(OH)3

 

 

(s)

 

 

Al2O3

 

 

(s)

 

 

+

 

 

3H2O(l)

 

 

 

12

 

o

Pure aluminium (III)oxide (Al2O3) has a very high melting point of 2015 C.

 

 

It is therefore dissolved first in molten cryolite /sodium hexafluoroaluminate (III)/Na3AlF6 to lower the melting point to about 800oC.

 

 

The molten electrolyte is put in the Hall cell made up of a steel tank lined with carbon graphite and an anode suspended into the electrolyte. During the electrolysis:

 

(i)At the cathode;

4Al3+(l)  +  12e  4Al(l)

 

(ii) At the anode;
6O2-(l) 3O2(g)+  12e

 

Aluminium is denser than the electrolyte therefore sink to the bottom of the Hall cell.

 

At this temperature ,the Oxygen evolved/produced at the anode reacts with carbon anode to form carbon(IV)oxide gas that escape to the atmosphere.

 

C(s)   + O2(g)    CO2(g)

The anode thus should be continuously replaced from time to time.

 

Flow chart summary of extraction of aluminium from Bauxite

 

 

 

 

 

 

 

 

Hot concentrated sodium

 

hydroxide

 

 

 

 

Bauxite(Al2O3.2H2O) ore with
Crush (increase Powdered mixture
impurities Fe2O3 and SiO2 surface area)

 

 

13

 

 

Carbon(IV)oxide

 

 

 

 

 

 

 

 

 

Aluminium hydroxide (Al(OH)3) as residue

 

 

 

 

 

 

Roast at

1000oC

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Pure aluminium

 

sinks in Hall cell

 

Sodium aluminate (NaAl(OH)4)

 

and sodium silicate (Na2SiO3)

 

as filtrate

 

 

 

 

 

Cryolite

 

 

 

 

 

 

 

Aluminium

 

  • Oxide

 

 

 

Electrolysis

 

 

Iron(III)oxide- Fe2O3 as residue

 

 

 

 

 

 

 

Sodium silicate  (Na2SiO3)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Oxygen  gas at anode

 

 

 

 

 

 

  1. Diagram showing the Hall cell / process for extraction of Bauxite

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

14

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

d)Uses of aluminium

 

  • In making aeroplane parts, buses, tankers, furniture because aluminium is very light.

 

(ii)Making duralumin-an alloy which is harder and has a higher tensile strength

 

(iii)Making utensils,sauce pans,spoons because it is light and good conductor of electricity.

 

(iv)Making overhead electric cables because it is light,ductile and good conductor of electricity.

 

(iv)Used in the thermite process for production of Manganese, Chromium amd Titanium.

 

 

  1. Environmental effects of extracting aluminium from Bauxite. Carbon(IV)oxide gas that escape to the atmosphere is a green house gas that causes global warming.

 

Bauxite is extracted by open caste mining that causes soil/environmental degradation.

 

15

 

  1. f) Test for presence of Al3+

If an ore is suspected to contain Al3+ it is;

 

(i)added hot concentrated sulphuric(VI)/Nitric(V)acid to free the ions present.

 

(ii)the free ions are then added a precipitating reagent like 2M sodium hydroxide /2M aqueous ammonia.

 

Observation Inference
White precipitate in excess 2M NaOH(aq) Pb2+ , Al3+,  Zn2+
White precipitate in excess 2M NH3(aq) Pb2+ , Al3+
No black precipitate on adding Na2S(aq) Al3+
No white precipitate on adding either Al3+
NaCl(aq),HCl(aq),H2SO4(aq),Na2SO4(aq)

 

Practice

 

1.An unknown rock X was discovered in Ukraine. Test with dilute sulphuric (VI)acid shows rapid effervescence with production of a colourless gas A that forms a white precipitate with lime water and colourless solution B. On adding 3cm3 of 2M sodium hydroxide, a white precipitate C is formed that dissolves to form a colourless solution D on adding more sodium hydroxide. On adding 2M aqueous ammonia, a white precipitate E is formed which persist in excess aqueous ammonia.On which on adding 5cm3 of 1M Lead(II)nitrate(V) to F a white precipitate G is formed which remains on heating.

 

Identify:

 

A

Hydrogen/H2

B

Aluminium sulphate(VI)/Al2(SO4) 3

C

Aluminium hydroxide/ Al(OH4) 3

D

Tetrahydroxoaluminate(III)/ [Al(OH4) 3]

E

Aluminium hydroxide/ Al(OH) 3

F

Aluminium chloride/ AlCl3

 

 

 

 

16

 

2.Aluminium is obtained from the ore with the formula Al2O3. 2H2O. The ore is first heated and refined to obtain pure aluminium oxide (Al2O3). The oxide is then electrolysed to get Aluminium and oxygen gas using carbon anodes and carbon as cathode. Give the common name of the ore from where aluminium is extracted from ½ mark

 

What would be the importance of heating the ore first before refining it?1 mark To remove the water of crystallization

 

The refined ore has to be dissolved in cryolite first before electrolysis. Why is this necessary? 1½ mark

 

To lower the melting point of aluminium oxide from about 2015oC to 900oC so as to lower /reduce cost of production

 

Why are the carbon anodes replaced every now and then in the cell for electrolysing

aluminium oxide?                                                                   1 mark

 

Oxygen produced at anode react with carbon to form carbon(IV)oxide gas that escape

 

State two uses of aluminium

 

In making aeroplane parts, buses, tankers, utensils, sauce pans,spoons Making overhead electric cables

 

Making duralumin

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

17

 

  1. IRON

a)Natural occurrence

 

Iron is the second most common naturally occurring metal. It makes 4% of the earths crust as:

(i)Haematite(Fe2O3)

(ii)Magnetite(Fe3O4)

(iii)Siderite(FeCO3)

b)The blast furnace for extraction of iron from Haematite and Magnetite

 

a)Raw materials:

(i)Haematite(Fe2O3)

(ii)Magnetite(Fe3O4)

(iii)Siderite(FeCO3)

(iv)Coke/charcoal/ carbon

(v)Limestone

 

b)Chemical processes:

 

Iron is usually extracted from Haematite (Fe2O3), Magnetite(Fe3O4) Siderite (FeCO3).These ores contain silicon(IV)oxide(SiO2) and aluminium(III)oxide (Al2O3) as impurities.

 

When extracted from siderite, the ore must first be roasted in air to decompose the iron(II)Carbonate to Iron(II)oxide with production of carbon(IV)oxide gas:

 

FeCO3(s)     FeO(s)           +          CO2(g)

 

Iron(II)oxide is then rapidly oxidized by air to iron(III)oxide(Haematite).

 

4FeO(s)   +             O2(g)  2Fe2O3(s)

 

Haematite (Fe2O3), Magnetite(Fe3O4), coke and limestone are all then fed from top into a tall (about 30metres in height) tapered steel chamber lined with refractory bricks called a blast furnace.

 

The furnace is covered with inverted double cap to prevent/reduce amount of any gases escaping .

 

Near the base/bottom, blast of hot air at about 1000K (827oC) is driven/forced into the furnace through small holes called Tuyeres.

 

As the air enters ,it reacts with coke/charcoal/carbon to form carbon(IV)oxide gas. This reaction is highly exothermic.

 

 

C(s)+ O2(g)

 

 

CO2

 

 

(g)

 

 

∆H =

 

 

-394kJ

 

 

 

18

 

This raises the temperature at the bottom of the furnace to about 2000K(1650oC).As Carbon(IV)oxide gas rises up the furnace it reacts with more coke to form carbon(II)oxide gas.This reaction is endothermic.

 

CO2 (g) +  C(s)               2CO (g) ∆H = +173kJ

 

Carbon(II)oxide gas is a strong reducing agent that reduces the ores at the upper parts of the furnace where temperatures are about 750K(500oC) i.e. For Haematite;

 

 

Fe2O3

 

(s)

 

+

 

3CO(g)

 

 

2Fe(s) + CO2(g)

 

 

For Magnetite;

 

 

Fe3O4

 

(s)

 

+

 

4CO(g)

 

 

3Fe(s) + 4CO2(g)

 

 

 

Iron is denser than iron ore. As it falls to the hotter base of the furnace it melts and can easily be tapped off.

 

Limestone fed into the furnace decomposes to quicklime/calcium oxide and produce more carbon(IV)oxide gas.

 

CaCO3(s)    CaO(s) + CO2(g)

 

Quicklime/calcium oxide reacts with the impurities silicon(IV)oxide(SiO2) and aluminium(III)oxide(Al2O3)in the ore to form calcium silicate and calcium aluminate.

 

CaO(s) + SiO2(s)    CaSiO3 (l)

CaO(s) + Al2O3(s)   Ca Al2O4 (l)

 

 

Calcium silicate and calcium aluminate mixture is called slag.Slag is denser than iron ore but less dense than iron therefore float on the pure iron. It is tapped at different levels to be tapped off for use in:

 

(i)tarmacing roads

(ii) cement manufacture

(iii)as building construction material

 

(c)Uses of Iron

 

Iron obtained from the blast furnace is hard and brittle. It is called Pig iron. It is remelted, added scrap steel then cooled. This iron is called cast iron. Iron is mainly used to make:

 

(i)gates ,pipes, engine blocks, rails, charcoal iron boxes,lamp posts because it is cheap.

 

 

19

 

(ii)nails, cutlery, scissors, sinks, vats, spanners,steel rods, and railway points from steel.

 

Steel is an alloy of iron with carbon, and/or Vanadium, Manganese, Tungsten, Nickel ,Chromium. It does not rust/corrode like iron.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

  1. Environmental effects of extracting Iron from Blast furnace

 

(i)Carbon(IV)oxide(CO2) gas is a green house gas that causes/increases global warming if allowed to escape/leak from the furnace. (ii)Carbon(II)oxide(CO)gas is a highly poisonous/toxic odourless gas that can kill on leakage.

 

It is preferentially absorbed by the haemoglobin in mammals instead of Oxygen to form a stable compound that reduce free hemoglobin in the blood.

  • Haematite (Fe2O3), Magnetite(Fe3O4) and Siderite (FeCO3) are extracted through quarrying /open cast mining that cause soil / environmental degradation .

 

20

 

  1. Test for the presence of Iron

 

Iron naturally exist in its compound as Fe2+ /Fe3+ If an ore is suspected to contain Fe2+ /Fe3+ it is;

 

(i)added hot concentrated sulphuric(VI)/Nitric(V)acid to free the ions present.

 

(ii)the free ions are then added a precipitating reagent like 2M sodium hydroxide /2M aqueous ammonia which forms;

 

  1. an insoluble green precipitate in excess of 2M sodium hydroxide /2M aqueous ammonia if Fe2+ ions are present.

 

  1. an insoluble brown precipitate in excess of 2M sodium hydroxide /2M aqueous ammonia if Fe2+ ions are present.

 

 

 

Observation Inference
green precipitate in excess 2M NaOH(aq) Fe2+
green precipitate in excess 2M NH3(aq) Fe2+
brown precipitate in excess 2M NaOH(aq) Fe3+
brown precipitate in excess 2M NH3(aq) Fe3+

 

Practice questions

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

21

 

4.COPPER

 

a)Natural occurrence

 

Copper is found as uncombined element/metal on the earths crust in Zambia, Tanzania, USA and Canada .The chief ores of copper are:

(i)Copper pyrites(CuFeS2)

(ii)Malachite(CuCO3.Cu(OH)2)

(iii)Cuprite(Cu2O)

b)Extraction of copper from copper pyrites.

 

Copper pyrites are first crushed into fine powder. The powdered ore is the added water and oil. The purpose of water is to dissolve hydrophilic substances/particle. The purpose of oil is to make cover copper ore particle so as to make it hydrophobic

 

Air is blown through the mixture. Air creates bubbles that stick around hydrophobic copper ore. The air bubbles raise through buoyancy small hydrophobic copper ore particles to the surface. A concentrated ore floats at the top as froth. This is called froth flotation. The concentrated ore is then skimmed off.The ore is then roasted in air to form copper(I)sulphide ,sulphur(IV)oxide and iron (II) oxide.

2CuFeS2(s) + 4O2(g)     Cu2S(s) + 3SO2(g)  + 2FeO(s)

 

Limestone (CaCO3) and silicon(IV)oxide (SiO2) are added and the mixture heated in absence of air.Silicon(IV)oxide (SiO2) reacts with iron (II) oxide to form Iron silicate which constitutes the slag and is removed.

FeO(s)   +       SiO2(s) FeSiO3(s)

 

The slag separates off from the copper(I)sulphide. Copper(I)sulphide is then heated in a regulated supply of air where some of it is converted to copper (I) oxide.

2Cu2S (s)  +  3O2(g)       2Cu2S(s)  +   2SO2(g)

 

The mixture then undergo self reduction in which copper(I)oxide is reduced by copper(I)sulphide to copper metal.

Cu2S (s)  +  2Cu2O (s)      6Cu (s)  + SO2(g)

 

The copper obtained has Iron, sulphur and traces of silver and gold as impurities.It is therefore about 97.5% pure. It is refined by electrolysis/electrolytic method.

 

During the electrolysis of refining copper, the impure copper is made the anode and a small pure strip is made the cathode. Electrode ionization takes place where:

 

At the anode;

 

22

 

Cu(s)   Cu2+ (aq)  +       2e

Note: Impure copper anode dissolves/erodes into solution and decreases in

 

size.

At the Cathode;

Cu2+ (aq)  +       2e    Cu(s)

 

Note: The copper ions in the electrolyte(CuSO4) are reduced and deposited as copper metal at the cathode. The copper obtained is 99.98% pure.

 

Valuable traces of silver and gold collect at the bottom of the electrolytic cell as sludge. It is used to finance the extraction of copper pyrites.

 

 

 

(c)Flow chart summary of extraction of copper from Copper pyrites

 

 

 

 

Water Oil
Copper pyrites(CuFeS2) ore with Crush (increase
Froth
impurities Fe2O3 and SiO2 surface area)
Excess  air Rocky impurities
Silicon(IV)
oxide Limestone 1st roasting chamber Concentration
chamber

 

Cu2S

 

Sulphur(IV)Oxide

 

Smelting furnace

 

Limited air

 

 

 

Calcium Iron Silicate
2nd roasting furnace
aluminate (FeSiO3)Slag Cu2S, Cu2O
(CaAl2O4)slag
Impure copper Self reduction

 

Sulphur(IV)Oxide

 

Electrolysis using Copper electrodes

 

 

23

 

Cathode; Pure Copper deposited.

Anode; Impure Copper eroded.

 

 

 

 

Electrolytic purification of impure copper

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

  1. Uses of copper Copper is mainly used in:

 

(i)making low voltage electric cables,contact switches, cockets and plugs because it is a good conductor of electricity.

 

(ii)Making solder because it is a good thermal conductor. (iii)Making useful alloys e.g.

 

-Brass is an alloy of copper and Zinc(Cu/Zn) -Bronze is an alloy of copper and Tin(Cu/Sn)

 

-German silver is an alloy of copper ,Zinc and Nickel(Cu/Zn/Ni) (iv)Making coins and ornaments.

 

 

  1. Environmental effects of extracting copper from Copper pyrites (i)Sulphur(IV)oxide is a gas that has a pungent poisonous smell that causes head ache to human in high concentration.

 

(ii)Sulphur(IV)oxide gas if allowed to escape dissolves in water /rivers/rain to form weak sulphuric(IV)acid lowering the pH of the water leading to marine pollution, accelerated corrosion/rusting of metals/roofs and breathing problems to human beings.

 

 

24

 

(iii)Copper is extracted by open caste mining leading to land /environmental /soil degradation.

 

  1. Test for the presence of copper in an ore Copper naturally exist in its compound as Cu2+ /Cu+ Copper (I) / Cu+ is readily oxidized to copper(II)/ Cu2+ If an ore is suspected to contain Cu2+ /Cu+ it is;

 

(i)added hot concentrated sulphuric(VI)/Nitric(V)acid to free the ions present.

 

(ii)the free ions are then added a precipitating reagent; 2M sodium hydroxide /2M aqueous ammonia which forms;

 

  1. an insoluble blue precipitate in excess of 2M sodium hydroxide if Cu2+ ions are present.

 

  1. an insoluble blue precipitate in 2M aqueous ammonia that dissolve to royal/deep blue solution in excess if Cu2+ ions are present.

 

 

 

Observation Inference
blue precipitate in excess 2M NaOH(aq) Cu2+
blue precipitate,dissolve to royal/deep blue Cu2+
solution in excess 2M NH3(aq)

 

g)Sample questions

 

Copper is extracted from copper pyrites as in the flow chart outlined below. Study it and answer the questions that follow

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

25

 

 

 

 

 

 

 

 

5.ZINC and LEAD

 

a)Natural occurrence

 

Zinc occurs mainly as:

(i)Calamine-Zinc carbonate(ZnCO3)

(ii)Zinc blende-Zinc sulphide(ZnS)

Lead occurs mainly as Galena-Lead(II)Sulphide mixed with Zinc blende:

 

b)Extraction of Zinc/Lead  from Calamine ,Zinc blende and Galena.

 

During extraction of Zinc , the ore is first roasted in air:

 

For Calamine Zinc carbonate decompose to Zinc oxide and carbon(IV) oxide gas.

ZnCO3(s)   ZnO(s)         +             CO2(g)

 

Zinc blende does not decompose but reacts with air to form Zinc oxide and sulphur(IV) oxide gas.

 

Galena as a useful impurity also reacts with air to form Lead(II) oxide and sulphur(IV) oxide gas.

 

 

2ZnS(s)  + 3O2(g)

 

 

2ZnO(s)

 

 

+

 

 

2SO2(g)

 

 

(Zinc blende)

 

 

2PbS(s)

 

 

+ 3O2(g)

 

 

2PbO(s)

 

 

+

 

 

2SO2(g)

 

 

(Galena)

 

 

The oxides are mixed with coke and limestone/Iron(II)oxide/ Aluminium (III) oxide and heated in a blast furnace.

 

At the furnace temperatures limestone decomposes to quicklime/CaO and produce Carbon(IV)oxide gas.

 

CaCO3(s)  CaO(s) + CO2 (g)

 

Carbon(IV)oxide gas reacts with more coke to form the Carbon(II)oxide gas.

C(s)  +  CO2 (g)    2CO (g)

 

Both Carbon(II)oxide and carbon/coke/carbon are reducing agents.

 

26

 

 

The oxides are reduced to the metals by either coke or carbon (II)oxide.

 

ZnO(s) + C(s) Zn(g)  + CO (g)
PbO(s) + C(s) Pb(l) + CO (g)
PbO(s) + CO(s) Pb(l) + CO2 (g)
PbO(s) + CO(s) Pb(g)  + CO2 (g)

 

At the furnace temperature:

 

(i)Zinc is a gas/vapour and is collected at the top of the furnace. It is condensed in a spray of molten lead to prevent reoxidation to Zinc oxide. On further cooling , Zinc collects on the surface from where it can be tapped off

 

(ii)Lead is a liquid and is ale to trickle to the bottom of the furnace from where it is tapped off.

 

Quicklime/CaO, Iron(II)Oxide, Aluminium(III)oxide are used to remove silica/silicon(IV)oxide as silicates which float above Lead preventing its reoxidation back to Lead(II)Oxide.

CaO(s) + SiO2(s)  CaSiO3(s/l)

(Slag-Calcium silicate)

FeO(s) + SiO2(s)  FeSiO3(s/l)

(Slag-Iron silicate)

Al2O3(s) + SiO2(s)  Al2SiO4(s/l)

(Slag-Aluminium silicate)

 

c)Flow chart on extraction of Zinc  from Calamine ,Zinc blende.

 

Oil Water CO2 from SO2 from
Zinc ore (calamine Powdered ore
calamine Zinc blende
/Zinc blende

 

 

Froth flotation Roasting chamber

 

 

 

Iron/aluminium/ Limestone

 

Reduction chamber
Coke

 

 

Water Filtration
Condenser
27 Slag (Iron silicate/
aluminium
silicate/calcium

 

 

 

 

 

 

 

 

 

 

  1. d) Flow chart on extraction of Lead from Galena

 

 

Lead ore/Galena Powdered ore oil Water
Iron/Limestone coke

Froth flotation

 

Roasting chamber

 

 

Reduction chamber

 

SO2(g)

 

 

 

Slag(Iron silicate)

 

 

Filtration
Condenser LEAD VAPOUR

 

 

 

Zinc residue

 

 

  1. e) Uses of Lead

 

Lead is used in:

(i)making gun-burettes.

 

(ii)making protective clothes against nuclear (alpha rays/particle) radiation in a nuclear reactor.

(iii)Mixed with tin(Sn) to make solder alloy

 

  1. Uses of Zinc Zinc is used in:

 

28

 

(i)Galvanization-when iron sheet is dipped in molten Zinc ,a thin layer of Zinc is formed on the surface.Since Zinc is more reactive than iron ,it reacts with elements of air(CO2/ O2 / H2O) to form basic Zinc carbonate(ZnCO3.Zn(OH)2).This sacrificial method protects iron from corrosion/rusting.

 

(ii)As negative terminal and casing in dry/Laclanche cells.

(iii)Making brass alloy with copper(Cu/Zn)

 

  1. Environmental effects of extracting Zinc and Lead.

 

(i) Lead and Lead salts are carcinogenic/causes cancer (ii)Carbon(IV)oxide is a green house gas that causes/accelerate global warming.

 

(iii)Carbon(II)oxide is a colourless odourless poisonous /toxic gas that combines with haemoglobin in the blood to form stable carboxyhaemoglobin reducing free haemoglobin leading to death.

 

(iv) Sulphur(IV)oxide is a gas that has a pungent poisonous smell that causes headache to human if in high concentration.

 

(v)Any leakages in Sulphur(IV)oxide gas escapes to the water bodies to form weak sulphuric(VI)acid lowering the pH of the water. This causes marine pollution /death of aquatic life, accelerated rusting/corrosion of metals/roofs and breathing problems to human beings.

 

  1. Test for presence of Zinc/ Lead.

 

If an ore is suspected to contain Zinc/Lead it is:

 

I.added hot concentrated Nitric(V)acid to free the ions present.

Note:

Concentrated Sulphuric(VI)acid forms insoluble PbSO4 thus cannot be used to free the ions in Lead salts.

 

II.the free ions are then added a precipitating reagent mostly 2M sodium hydroxide or 2M aqueous ammonia with the formation of;

 

(i)a soluble precipitate in excess of 2M sodium hydroxide if Zn2+, Pb2+, Al3+ions are present.

 

(ii)a white precipitate that dissolves to form a colorless solution in excess 2M aqueous ammonia if Zn2+ions are present.

 

(iii)an insoluble white precipitate in excess 2M aqueous ammonia if Pb2+, Al3+ions are present.

 

  • Pb2+ ions form a white precipitate when any soluble SO42-, SO32-, CO32-, Cl, is added while Al3+ ions do not form a white precipitate

 

29

 

  • Pb2+ ions form a yellow precipitate when any soluble I(e.g. Potassium/sodium Iodide) is added while Al3+ ions do not form a yellow precipitate.

 

  • Pb2+ ions form a black precipitate when any soluble S(e.g. Potassium/sodium sulphide) is added while Al3+ ions do not form a black precipitate.i.e;

 

Observation Inference
White precipitate in excess 2M NaOH (aq) Zn2+, Pb2+, Al3+ ions
White precipitate that dissolves to form a Zn2+ ions
colourless solution in excess 2M NH3(aq)
White precipitate in excess 2M NH3(aq) Pb2+, Al3+ ions
White precipitate on adding about 4 drops of Pb2+ions
either  Na2CO3(aq), Na2SO3(aq), Na2SO4(aq),
H2SO4(aq), HCl(aq), NaCl(aq)
Yellow precipitate on adding about 4 drops of Pb2+ ions
of KI(aq).NaI (aq)
Black precipitate on adding aout 4 drops of Pb2+ ions
Na2S(aq)/K2S(aq)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

30

 

 

6.GENERAL SUMMARY OF METALS

 

  1. a) Summary methods of extracting metal from their ore

 

The main criteria used in extraction of metals is based on its position in the electrochemical/reactivity series and its occurrence on the earths crust.

 

Position on the earth’s crust

 

 

 

 

 

 

 

 

If near the surface use open cast mining / quarrying

 

 

 

 

 

 

 

If deep on the earth’s

 

crust use deep mining

 

 

 

 

 

 

Add oil, water, and blow air to form froth

 

to concentrate the ore if it is a low grade

 

 

 

 

 

 

 

 

 

 

 

Electrolyse the ore if it is made of reactive metals; Potassium, Sodium, Magnesium, Calcium, Aluminium

 

 

 

 

 

 

 

 

 

 

Roast the ore first if it is a carbonate / sulphide of Zinc, Iron, Tin, Lead, and copper to form the oxide

 

 

 

 

 

 

 

Reduce the oxide using carbon in a furnace if it is made of Zinc ,Tin, Lead ,Copper and Iron

 

 

 

 

31

 

 

 

 

 

 

 

  1. b) Summary of extraction of common metal.
Metal Chief Chemical Method of extraction Main equation during
ore/s formula of extraction
ore
Sodium Rock salt NaCl(s) Downs process Cathode:
Through electrolysis of 2Na+(l) + 2e -> 2Na(l)
molten NaCl (CaCl2 Anode:
lower m.pt from 800oC-> 2Cl(l) -> Cl2(g) + 2e
600oC)
Sodium/ Brine NaCl(aq) Flowing mercury Cathode:
sodium cathode cell 2Na+(aq)+2e ->2Na(aq)
hydroxide Through electrolysis of Anode:
concentrated NaCl(aq) 2Cl(aq) -> Cl2(g) + 2e
Aluminium Bauxite Al2O3.2H2O Halls process Cathode:
Through electrolysis of 4Al3+(l) + 12e -> 4Al(l)
molten Al2O3. (Cryolite Anode:
lower m.pt from 2015oC 6O2-(l) -> 3O2(g) + 12e
-> 800oC)
Iron Haematite Fe2O3 Blast furnace Fe2O3(s)+ 3CO(g)
Magnetite Fe3O4 Reduction of the ore by 2Fe(l) +3CO2(g)
carbon(II)oxide
Fe3O4(s)+ 4CO(g)
3Fe(l) +4CO2(g)
Copper Copper CuFeS2 Roasting the ore in air 2CuFeS2 (s)+ 4O2(g) ->
pyrites to get Cu2S. Cu2S(s)+3SO2(g)
Heating Cu2S ore in +2FeO(s)
regulated supply of air.
Reduction of Cu2O  by 2Cu2S (s)+ 3O2(g) ->
Cu2S 2Cu2O(s)+2SO2(g)
Cu2S (s)+ 2Cu2O(s) ->
6Cu(s)+ SO2(g)
Zinc Calamine ZnCO3 Roasting the ore in air ZnCO3(s)-> ZnO(s) +
to get ZnO CO2(g)
32

 

Blast furnace
/reduction of the oxide 2ZnS(s) +3O2(g) ->
by 2ZnO(s) + 2SO2(g)
Carbon(II)Oxide/Carbon
ZnO(s) + CO(g)->
Zn(s) + CO2(g)
Lead Galena PbS Blast furnace- PbO(s) + CO(g)->
Reduction of the oxide Pb(s) + CO2(g)
by carbon(II)oxide
/Carbon

 

  1. c) Common alloys of metal.
Alloy name Constituents of the alloy Uses of the alloy
Brass Copper and Zinc Making scews and bulb caps
Bronze Copper and Tin Making clock springs,electrical contacts
and copper coins
Soldier Lead and Tin Soldering, joining electrical contacts
because of its low melting points and high
thermal conductivity
Duralumin Aluminium, Copper and Making aircraft , utensils ,windows frames
Magnesium because of its light weight and corrosion
resistant.
Steel Iron, Carbon ,Manganese Railway lines , car bodies girders  and
and other metals utensils.
Nichrome Nichrome and Chromium Provide resistance in electric heaters and
ovens
German silver Copper,Zinc and Nickel Making coins

 

 

 

  1. d) Physical properties of metal.

 

Metals form giant metallic structure joined by metallic bond from electrostatic attraction between the metallic cation and free delocalized electrons. This makes metals to have the following physical properties:

 

(i)High melting and boiling points

 

The giant metallic structure has a very close packed metallic lattice joined by strong electrostatic attraction between the metallic cation and free delocalized electrons.The more delocalized electrons the higher the melting/boiling points e.g.

 

Aluminium has a melting point of about 2015oC while that of sodium is about 98oC.This is mainly because aluminium has more/three delocalized electrons than sodium/has one. Aluminium has a boiling point of about 2470oC while that of sodium is about 890oC.This is mainly because aluminium has more/three delocalized electrons than sodium/has one.

 

33

 

(ii)High thermal and electrical conductivity

 

All metals are good thermal and electrical conductors as liquid or solids. The more delocalized electrons the higher the thermal and electrical conductivity. e.g. Aluminium has an electrical conductivity of about 3.82 x 19-9 ohms per metre. Sodium has an electrical conductivity of about 2.18 x 19-9 ohms per metre.

 

(iii)Shiny/Lustrous

The free delocalized electrons on the surface of the metal absorb, vibrate and then

scatter/re-emit/lose light energy. All metals are therefore usually shades of grey in colour

except copper which is shiny brown.e.g.

 

Zinc is bluish grey while iron is silvery grey.

 

(iv)High tensile strength

 

The free delocalized electrons on the surface of the metal atoms binds the surface immediately when the metal is coiled/folded preventing it from breaking /being brittle.

 

(v)Malleable.

 

Metals can be made into thin sheet. The metallic crystal lattice on being beaten/pressed/hammered on two sides extend its length and width/bredth and is then immediately bound by the delocalized electrons preventing it from breaking/being brittle.

 

(vi)Ductile.

 

Metals can be made into thin wires. The metallic crystal lattice on being beaten/pressed/hammered on all sides extend its length is then immediately bound by the delocalized electrons preventing it from breaking/being brittle.

 

Revision  questions

 

1.Given some soil , dilute sulphuric(VI)acid,mortar,pestle,filter paper,filter funnel and 2M aqueous ammonia,describe with explanation,how you would show that the soil contain Zinc.

 

Place the soil sample in the pestle. Crush using the mortar to reduce the particle size/increase its surface area.

Add dilute sulphuric(VI)acid to free the ions in soil sample.

Filter to separate insoluble residue from soluble filtrate

 

To filtrate,add three drops of aqueous ammonia as precipitating reagent. A white precipitate of Zn(OH)2, Pb(OH)2 or Al(OH)3 is formed

Add excess aqueous ammonia to the white precipitate. If it dissolves the Zn2+ ions are present. Zn(OH)2 react with excess ammonia to form soluble [Zn(OH)4]2+ complex.

 

2.In the extraction of aluminium,the oxide is dissolved in cryolite.

(i)What is the chemical name of cryolite?

Sodium hexafloroaluminate/Na3AlF6

 

 

 

34

 

(ii)What is the purpose of cryolite?

 

To lower the melting point of the electrolyte/Aluminium oxide from about 2015oC to 900oC

 

(iii)Name the substance used for similar purpose in the Down cell Calcium chloride/CaCl2

 

(iv)An alloy of sodium and potassium is used as coolant in nuclear reactors.Explain. Nuclear reactors generate a lot of heat energy. sodium and potassium alloy reduce/lower the high temperature in the reactors.

 

(v)Aluminium metal is used to make cooking utensils in preference to other metals.Explain.

 

Aluminium

 

  • is a very good conductor of electricity because it has three delocalized electrons in its metallic structure

 

(ii)is cheap,malleable,ductile and has high tensile strength

 

(iii)on exposure to fire/heat form an impervious layer that prevent it from rapid corrosion.

 

3.Study the scheme below and use it to answer the questions that follow.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(a)Identify:

 

(i)solid residue L

Iron(III)Oxide/Fe2O3

 

(ii)Solid N

Aluminium hydroxide /Al(OH)3

 

 

35

 

(iii)Filtrate M

Sodium tetrahydroxoaluminate/ NaAl(OH)4 and sodium silicate/ NaSiO3

 

(iv)Solid P

Aluminium oxide/ Al2O3

 

(v)Gas Q

Oxygen/O2

 

(vi)Process K1

Filtration

 

(vii)Process K2

Electrolysis

 

(b)Write the equation for the reaction taking place in the formation of solid P from solid N

2Al(OH)3      ->          Al2O3 (s)       +            3H2O(l)

 

(c)Name a substance added to solid N before process Process K2 take place. Cryolite/Sodium tetrahydroxoaluminate/ NaAl(OH)4

 

(d)State the effect of evolution of gas Q on (i)process K2

 

Oxygen produced at the anode reacts with the carbon anode to form carbon(IV) oxide which escape. The electrolytic process needs continuous replacement of the carbon anode.

 

(ii)the environment

 

Oxygen produced at the anode reacts with the carbon anode to form carbon(IV) oxide which escape to the atmosphere.CO2 is a green house gas that cause global warming.

 

(e)An aluminium manufacturing factory runs for 24 hours. If the total mass of aluminium produced is 27000kg,

 

(i)Calculate the current used. (Faraday constant=96500Coulombs, Al=27.0). (ii)assuming all the gas produced react with 200kg of anode ,calculate the loss in mass of the electrode.(Molar gas volume at room temperature = 24dm3,C=12.0)

 

Working

Equation at Cathode

Al3+(l)              +            3e          ->          Al(l)

27g  Al   ->  3 Faradays   =   3   x     96500C

 

(27000kg x 1000) g -> (27000kg x 1000) g x 3 x 96500C 27g

 

=289500000000 Coulombs

 

36

 

Current  = Quantity of electricity   =>289500000000 Coulombs
Time in seconds 24 x 60 x 60
3350690Ampheres
Working
Equation at Anode
2O2-(l) + 4e -> O2(g)
4 Faradays ->  4  x  96500C24dm3 O2(g)  –
289500000000 Coulombs  ->  289500000000 Coulombs x 24dm3
4 x  96500C
18,000,000dm3
Chemical equation at anode
O2(g) + C (s) -> CO2(g)
Method 1
24dm3 of O2(g) -> 12.0g Carbon
18,000,000dm3 of O2(g) -> 18,000,000dm3 x 12 = 9000000g = 9000kg
24dm3 1000g

 

Loss in mass of the carbon graphite anode =  9000kg

NB:Mass of the carbon graphite anode  remaining =27000kg –  9000kg =18000kg

 

The flow chart below shows the extraction of iron metal.Use it to answer the questions that follow.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(a)Identify:

(i)gas P

Carbon(IV)oxide/CO2

 

(ii)Solid Q

 

37

 

Carbon/coke/charcoal

 

(iii)Solid R

Carbon/coke/charcoal

 

(iv)Solid V

Limestone/calcium carbonate/CaCO3

 

(v)Solid S

Iron/Fe

 

(b)Write the chemical equation for the reaction for the formation of:

 

(i)Solid S

Fe2O3(s)  +   3CO(g)   ->   2Fe(s)   +    3CO2(g)

 

(ii)Carbon(II)oxide

C(s)   +   CO2 (g)  ->   2CO (g)

 

 

(iii)Slag SiO2(s) Al2O3 (s)

 

 

 

 

+   CaO(s)   ->   CaSiO3(s)

+   CaO(s)   ->   Ca Al2O4(s)

 

 

 

(iv)Gas P

C(s)

 

 

 

+    O2 (g)  ->   CO2 (g)

 

 

(c)State two uses of:

 

(i)Solid S

Iron is used in making:

 

(i)gates ,pipes, engine blocks, rails, charcoal iron boxes, lamp posts because it is cheap. (ii)nails, cutlery, scissors, sinks, vats, spanners, steel rods, and railway points from steel. Steel is an alloy of iron with carbon, and/or Vanadium, Manganese, Tungsten, Nickel ,Chromium.

 

It does not rust/corrode like iron.

 

(ii)Slag

(i) tarmacing roads

  • cement manufacture
  • as building construction material

 

3.You are provided with sulphuric(VI)acid ,2M aqueous ammonia and two ores suspected to contain copper and iron. Describe with explanation how you would differentiate the two ores.

 

Crush the two ores separately in using a mortar and pestle to reduce the particle size/increase the surface area.

Add sulphuric(VI)acid  to separate portion of the ore. Filter.

 

38

 

To a portion of the filtrate,add three drops of 2M aqueous ammonia then axcess Results

 

A green precipitate insoluble in excess 2M aqueous ammonia confirms the ore contain Fe2+ ion.

 

A brown precipitate insoluble in excess 2M aqueous ammonia confirms the ore contain Fe3+ ion.

 

A blue precipitate that dissolve in excess 2M aqueous ammonia to form a deep/royal blue solution confirms the ore contain Cu2+ ion.

 

  1. Use the flow chart below showing the extraction of Zinc metal to answer the questions that follow

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

(a)Name:

 

(i)two ores from which Zinc can be extracted Calamine(ZnCO3)

Zinc blende(ZnS)

 

(ii)two possible identity of gas P

Sulphur(IV)oxide(SO2) from roasting Zinc blende

Carbon(IV)oxide(CO2) from  decomposition of Calamine.

 

(b)Write a possible chemical equation taking place in the roasting chamber.

2ZnS(s)           +            3O2 (g) ->    2ZnO(s)          +   2SO2(g)

ZnCO3(s)         ->          ZnO(s)           +   CO2(g)

 

(c)Explain the effect of the by-product of the roating on the environment.

Sulphur (IV)oxide from roasting Zinc blende is an acidic gas that causes acid rain on dissolving in rain water.

 

Carbon(IV)oxide(CO2) from decomposition of Calamine is a green house gas that causes global warming.

 

(d)(i)Name a suitable reducing agent used in the furnace during extraction of  Zinc.

Carbon(II)oxide

 

 

39

 

(ii)Write a chemical equation  for the reduction process

ZnO(s)    +            CO(g) ->          Zn(s)                +    CO2(g)

 

(e)(i)Before electrolysis, the products from roasting is added dilute sulphuric (VI)acid. Write the equation for the reaction with dilute sulphuric(VI)acid.

ZnO(s) +  H2SO4 (aq)  ->  Zn SO4(aq) +  H2(g)

 

(ii)During the electrolysis for extraction of Zinc,state the I. Anode used

Aluminium sheet

  1. Cathode used

 

Lead plate coated with silver

(ii)Write the equation for  the electrolysis for extraction of Zinc at the:

I.Cathode;

 

Zn2+(aq)

 

+

 

2e   ->     Zn(s)

 

 

 

II.Anode;

 

 

4OH-(aq)

 

->  2H2O(l)  +  O2(s) +

 

4e

 

 

 

(f)(i)What is galvanization

Dipping Iron in molten Zinc to form a thin layer of Zinc to prevent iron  from rusting.

 

(ii)Galvanized iron sheet rust after some time. Explain

 

The thin layer of Zinc protect Iron from rusting through sacrificial protection. When all the Zinc has reacted with elements of air, Iron start rusting.

 

(g)State two uses of Zinc other than galvanization.

Making brass(Zinc/copper alloy)

Making  german silver(Zinc/copper/nickel alloy)

As casing for dry cells/battery

 

(h)Calculate the mass of Zinc that is produced from the reduction chamber if 6400kg of Calamine ore is fed into the roaster. Assume the process is 80% efficient in each stage(Zn=64.0,C=12.0,O=16.0)

 

Molar mass ZnCO3(s) =124g
Molar mass Zn = 64g
Molar mass ZnO = 80g
Chemical equation
ZnCO3(s)->ZnO(s) +  CO2(g)
Method 1
124g  ZnCO3 => 80g  ZnO

 

40

 

(6400kg x1000)g ZnCO3 => (6400  x1000) x 80 = 512,000,000 g of ZnO
124
100% => 512,000,000 g of ZnO
80% => 80 x  512,000,000 g 409600000g of ZnO
100
Chemical equation
ZnO(s) + CO(g) -> Zn(s) + CO2(g)
80g ZnO(s) => 64g Zn(s)
409600000g of ZnO => 409600000g x  64 =   327680000 g Zn
80
100% => 327680000 g Zn
80% => 80 x 327680000 g Zn = 262144000g of Zn
100
Mass of Zinc produced  =  262144000g  of  Zn

 

5.An ore is suspected to bauxite. Describe the process that can be used to confirm the presence of aluminium in the ore.

 

Crush the ore to fine powder to increase surface area/reduce particle size.

Add hot concentrated sulphuric(VI)/nitric(V) acid to free the ions.

Filter. Retain the filtrate

Add excess aqueous ammonia to a sample of filtrate.

A white precipitate confirms presence of either Al3+ or Pb2+.

 

Add sodium sulphate,dilute sulphuric(VI)to another portion of filtrate. No white precipitate confirms presence of Al3+

Or Add potassium iodide to another portion of filtrate.

No yellow precipitate confirms presence of Al3+

 

6.The flow chart below illustrate the industrial extraction of Lead metal

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

41

 

(a)(i)Name the chief ore that is commonly used in this process Galena(PbS)

 

(ii)Explain what take place in the roasting furnace

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

42

 

 

23.0.0 ORGANIC CHEMISTRY II (ALKANOLS AND ALKANOIC ACIDS) (20 LESSONS)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

B.ALKANOLS(Alcohols)

 

(A) INTRODUCTION.

 

Alkanols belong to a homologous series of organic compounds with a general formula CnH2n +1 OH and thus -OH as the functional group .The 1st ten alkanols include

 

n General / Structural formula IUPAC
molecular name
formular
1 CH3OH HCO-H
Methanol
H
2 CH3 CH2OH H    H Ethanol

 

C2H5 OH C CO-H
H
H H
3 CH3 (CH2)2OH H H H Propanol
C3H7 OH
H C C -CO-H
H H H
4 CH3 (CH2)3OH H H H H Butanol
C4H9 OH
H C C- C- CO-H
H H H H
5 CH3(CH2)4OH H H H  H  H Pentanol
C5H11 OH
H C C- C-C- CO-H
H H H  H  H
6 CH3(CH2)5OH H H H  H  H H Hexanol
C6H13 OH
H C C- C-C- C C-O-H
H HHHHH
7 CH3(CH2)6OH H H H  H  H H H Heptanol
C7H15 OH
H C C- C-C- C CC-O-H
H HHHHH H
8 CH3(CH2)7OH H H H  H  H H H  H Octanol
C8H17 OH
H C C- C-C- C CC-C-O-H
H H H  H  H H  H  H
2

 

9 CH3(CH2)8OH H H H  H  H H H  H  H Nonanol
C9H19 OH
H C C- C-C- C CC-CC-O-H
H H H  H  H H H  H  H
10 CH3(CH2)9OH H H H  H  H H HHHH Decanol
C10H21 OH
H C C- C-C- C CC-CC-C-O-H
H H H  H  H H HH HH

 

Alkanols like Hydrocarbons( alkanes/alkenes/alkynes) form a homologous series where:

 

(i)general name is derived from the alkane name then ending with “-ol” (ii)the members have OH as the fuctional group

 

(iii)they have the same general formula represented by R-OH where R is an alkyl group.

(iv) each member differ by CH2 group from the next/previous.

 

(v)they show a similar and gradual change in their physical properties e.g. boiling and melting points.

 

(vi)they show similar and gradual change in their chemical properties.

 

  1. ISOMERS OF ALKANOLS.

 

Alkanols exhibit both structural and position isomerism. The isomers are named by using the following basic guidelines:

 

(i)Like alkanes , identify the longest carbon chain to be the parent name. (ii)Identify the position of the -OH functional group to give it the smallest /lowest position.

(iii) Identify the type and position of the side branches.

 

Practice examples of isomers of alkanols

 

(i)Isomers of propanol C3H7OH

CH3CH2CH2OH  – Propan-1-ol

OH

 

CH3CHCH3  – Propan-2-ol

 

 

3

 

Propan-2-ol and Propan-1-ol are position isomers because only the position of the OH functional group changes.

 

(ii)Isomers of  Butanol     C4H9OH

 

CH3 CH2 CH3 CH2 OH Butan-1-ol

 

CH3 CH2 CH CH3

 

 

OH                             Butan-2-ol

 

CH3

 

 

CH3 CH3 CH3

 

 

OH                                       2-methylpropan-2-ol

 

Butan-2-ol and Butan-1-ol are position isomers because only the position of the -OH functional group changes.

 

2-methylpropan-2-ol is both a structural and position isomers because both the position of the functional group and the arrangement of the atoms in the molecule changes.

 

(iii)Isomers of  Pentanol     C5H11OH

 

CH3 CH2 CH2CH2CH2 OH Pentan-1-ol (Position isomer)

 

CH3 CH2 CH CH3

 

 

OH                             Pentan-2-ol (Position isomer)

 

CH3 CH2 CH CH2 CH3

 

 

OH                             Pentan-3-ol (Position isomer)

 

 

 

 

CH3

 

 

CH3 CH2 CH2 C CH3

 

 

 

4

 

OH                               2-methylbutan-2-ol (Position /structural isomer)

 

CH3

 

 

CH3 CH2 CH2 C CHOH

 

 

CH3                                                                                                              2,2-dimethylbutan-1-ol (Position /structural isomer)

 

CH3

 

 

CH3 CH2 CH   C CH3

 

 

CH3  OH                2,3-dimethylbutan-1-ol (Position /structural isomer)

 

(iv)1,2-dichloropropan-2-ol
CClH2 CCl CH3
OH
(v)1,2-dichloropropan-1-ol
CClH2 CHCl CH2
OH
(vi) Ethan1,2-diol
H  H
HOCH2CH2OH H-O-C-CO-H
H   H
(vii) Propan1,2,3-triol H  OHH
HOCH2CHOHCH2OH H-O-C-CCO-H
H  H   H
  1. LABORATORY PREPARATION OF ALKANOLS.

 

For decades the world over, people have been fermenting grapes juice, sugar, carbohydrates and starch to produce ethanol as a social drug for relaxation.

 

5

 

 

In large amount, drinking of ethanol by mammals /human beings causes mental and physical lack of coordination.

 

Prolonged intake of ethanol causes permanent mental and physical lack of coordination because it damages vital organs like the liver.

 

Fermentation is the reaction where sugar is converted to alcohol/alkanol using

biological catalyst/enzymes in yeast.

It involves three processes:

(i)Conversion of starch to maltose using the enzyme diastase.

 

 

(C6H10O5)n (s) + H2O(l)

 

 

–diastase enzyme –> C12H22O11(aq)

 

 

(Starch)

 

(Maltose)

 

 

 

(ii)Hydrolysis of Maltose to glucose using the enzyme maltase.

 

 

C12H22O11(aq)+ H2O(l)

 

 

— maltase enzyme –>2 C6H12O6(aq)

 

 

(Maltose)

 

(glucose)

 

 

 

(iii)Conversion of glucose to ethanol and carbon(IV)oxide gas using the enzyme zymase.

 

 

C6H12O6(aq)

(glucose)

 

 

— zymase enzyme –> 2 C2H5OH(aq)

(Ethanol)

 

 

+ 2CO2(g)

 

 

 

At concentration greater than 15% by volume, the ethanol produced kills the yeast enzyme stopping the reaction.

To increases the concentration, fractional distillation is done to produce spirits (e.g.

Brandy=40% ethanol).

Methanol is much more poisonous /toxic than ethanol.

 

Taken large quantity in small quantity it causes instant blindness and liver, killing the consumer victim within hours.

 

School laboratory preparation of ethanol from fermentation of glucose

Measure 100cm3 of pure water into a conical flask.

Add about five spatula end full of glucose.

Stir the mixture to dissolve.

Add about one spatula end full of yeast.

 

Set up the apparatus as below.

 

 

 

 

 

 

6

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Preserve the mixture for about three days.

 

D.PHYSICAL AND CHEMICAL PROPERTIES OF ALKANOLS

 

Use the prepared sample above for the following experiments that shows the characteristic properties of alkanols

 

  • Role of yeast

 

Yeast is a single cell fungus which contains the enzyme maltase and zymase that catalyse the fermentation process.

 

  • Observations in lime water.

 

A white precipitate is formed that dissolve to a colourless solution later. Lime water/Calcium hydroxide reacts with carbon(IV)0xide produced during the fermentation to form insoluble calcium carbonate and water.

 

More carbon (IV)0xide produced during fermentation react with the insoluble calcium carbonate and water to form soluble calcium hydrogen carbonate.

 

 

Ca(OH)2(aq) + H2O(l) + CO2 (g)

 

 

 

 

+

 

 

CO2 (g) CaCO3(s) ->

 

 

->

 

 

CaCO3(s)

 

Ca(HCO3) 2 (aq)

 

 

 

(c)Effects on litmus paper

Experiment

 

7

 

Take the prepared sample and test with both blue and red litmus papers.

Repeat the same with pure ethanol and methylated spirit.

Sample Observation table

 

Substance/alkanol Effect on litmus paper
Prepared sample Blue litmus paper remain blue
Red litmus paper remain red
Absolute ethanol Blue litmus paper remain blue
Red litmus paper remain red
Methylated spirit Blue litmus paper remain blue
Red litmus paper remain red

Explanation

 

Alkanols are neutral compounds/solution that have characteristic sweet smell and taste.

They have no effect on both blue and red litmus papers.

 

(d)Solubility in water.

Experiment

 

Place about 5cm3 of prepared sample into a clean test tube Add equal amount of distilled water.

Repeat the same with pure ethanol and methylated spirit.

Observation

No layers formed between the two liquids.

Explanation

 

Ethanol is miscible in water.Both ethanol and water are polar compounds .

 

The solubility of alkanols decrease with increase in the alkyl chain/molecular mass. The alkyl group is insoluble in water while OH functional group is soluble in water.

 

As the molecular chain becomes longer ,the effect of the alkyl group increases as the effect of the functional group decreases.

 

e)Melting/boiling point.

Experiment

Place pure ethanol in a long boiling tube .Determine its boiling point.

Observation

 

Pure ethanol has a boiling point of 78oC at sea level/one atmosphere pressure. Explanation

 

The melting and boiling point of alkanols increase with increase in molecular chain/mass .

 

 

8

 

 

This is because the intermolecular/van-der-waals forces of attraction between the molecules increase.

 

More heat energy is thus required to weaken the longer chain during melting and break during boiling.

 

f)Density

 

Density of alkanols increase with increase in the intermolecular/van-der-waals forces of attraction between the molecule, making it very close to each other.

 

This reduces the volume occupied by the molecule and thus increase the their mass per unit volume (density).

 

Summary table showing the trend in physical properties of alkanols

 

Alkanol Melting point Boiling point Density Solubility in water
(oC) (oC) gcm-3
Methanol -98 65 0.791 soluble
Ethanol -117 78 0.789 soluble
Propanol -103 97 0.803 soluble
Butanol -89 117 0.810 Slightly soluble
Pentanol -78 138 0.814 Slightly soluble
Hexanol -52 157 0.815 Slightly  soluble
Heptanol -34 176 0.822 Slightly soluble
Octanol -15 195 0.824 Slightly soluble
Nonanol -7 212 0.827 Slightly soluble
Decanol 6 228 0.827 Slightly soluble

 

g)Burning

 

Experiment

 

Place the prepared sample in a watch glass. Ignite. Repeat with pure ethanol and methylated spirit.

 

Observation/Explanation

 

Fermentation produce ethanol with a lot of water(about a ratio of 1:3)which prevent the alcohol from igniting.

Pure ethanol and methylated spirit easily catch fire / highly flammable.

 

They burn with an almost colourless non-sooty/non-smoky blue flame to form carbon(IV) oxide (in excess air/oxygen)or carbon(II) oxide (limited air) and water.

 

Ethanol is thus a saturated compound like alkanes.

 

 

9

 

Chemica equation
C2 H5OH(l) + 3O2 (g) -> 3H2O(l) + 2CO2 (g)  ( excess air)
C2 H5OH(l) + 2O2 (g) -> 3H2O(l) + 2CO (g)  ( limited air)
2CH3OH(l) + 3O2 (g) -> 4H2O(l) + 2CO2 (g)  ( excess air)
2 CH3OH(l) + 2O2 (g) -> 4H2O(l) + 2CO (g)  ( limited air)
2C3 H7OH(l) + 9O2 (g) -> 8H2O(l) + 6CO2 (g)  ( excess air)
C3 H7OH(l) + 3O2 (g) -> 4H2O(l) + 3CO (g)  ( limited air)
2C4 H9OH(l) + 13O2 (g) -> 20H2O(l)   + 8CO2 (g)  ( excess air)
C4 H9OH(l) + 3O2 (g) -> 4H2O(l) + 3CO (g)  ( limited air)

 

Due to its flammability, ethanol is used;

  • as a fuel in spirit lamps
  • as gasohol when blended with gasoline

 

(h)Formation of alkoxides

Experiment

 

Cut a very small piece of sodium. Put it in a beaker containing about 20cm3 of the prepared sample in a beaker.

 

Test the products with litmus papers. Repeat with absolute ethanol and methylated spirit.

 

Sample observations

Substance/alkanol Effect of adding sodium
Fermentation prepared sample (i)effervescence/fizzing/bubbles
(ii)colourless gas produced that
extinguish burning splint with explosion/
Popsound
(iii)colourless solution formed
(iv)blue litmus papers remain blue
(v)red litmus papers turn blue
Pure/absolute ethanol/methylated spirit (i)slow effervescence/fizzing/bubbles
(ii)colourless gas slowly  produced that
extinguish burning splint with explosion/
Popsound
(iii)colourless solution formed
(iv)blue litmus papers remain blue
(v)red litmus papers turn blue
Explanations
10

 

 

Sodium/potassium reacts slowly with alkanols to form basic solution called alkoxides and producing hydrogen gas.

 

If the alkanol has some water the metals react faster with the water to form soluble hydroxides/alkalis i.e.

 

Sodium + Alkanol -> Sodium alkoxides + Hydrogen gas
Potassium + Alkanol -> Potassium alkoxides + Hydrogen gas
Sodium + Water -> Sodium hydroxides + Hydrogen gas
Potassium + Water -> Potassium hydroxides + Hydrogen gas
Examples
1.Sodium metal reacts with ethanol to form sodium ethoxide
Sodium metal reacts with water to form sodium Hydroxide
2CH3CH2OH(l) + 2Na(s) -> 2CH3CH2ONa (aq) + H2 (s)
2H2O(l) + 2Na(s) -> 2NaOH (aq) + H2 (s)
2.Potassium metal reacts with ethanol to form Potassium ethoxide
Potassium metal reacts with water to form Potassium Hydroxide
2CH3CH2OH(l) + 2K(s) -> 2CH3CH2OK (aq) + H2 (s)
2H2O(l) + 2K(s) -> 2KOH (aq) + H2 (s)
3.Sodium metal reacts with propanol to form sodium propoxide
Sodium metal reacts with water to form sodium Hydroxide
2CH3CH2 CH2OH(l) + 2Na(s) -> 2CH3CH2 CH2ONa (aq)   + H2 (s)
2H2O(l) + 2Na(s) -> 2NaOH (aq) + H2 (s)
4.Potassium metal reacts with propanol to form Potassium propoxide
Potassium metal reacts with water to form Potassium Hydroxide
2CH3CH2 CH2OH(l) + 2K(s) -> 2CH3CH2 CH2OK (aq)   + H2 (s)
2H2O(l) + 2K(s) -> 2KOH (aq) + H2 (s)

 

5.Sodium metal reacts with butanol to form sodium butoxide Sodium metal reacts with water to form sodium Hydroxide

2CH3CH2 CH2 CH2OH(l)  + 2Na(s) -> 2CH3CH2 CH2 CH2ONa (aq) + H2 (s)

2H2O(l)          +              2Na(s)              ->          2NaOH (aq)     + H2 (s)

 

6.Sodium metal reacts with pentanol to form sodium pentoxide Sodium metal reacts with water to form sodium Hydroxide

2CH3CH2 CH2 CH2 CH2OH(l)+2Na(s) -> 2CH3CH2 CH2 CH2 CH2ONa (aq) + H2 (s)

 

11

 

2H2O(l)          +              2Na(s)              ->          2NaOH (aq)     + H2 (s)

 

(i)Formation of Esters/Esterification

Experiment

 

Place 2cm3 of ethanol in a boiling tube.

Add equal amount of ethanoic acid.To the mixture add carefully 2drops of

concentrated sulphuric(VI)acid.

Warm/Heat gently.

Pour the mixture into a beaker containing about 50cm3 of cold water.

Smell the products.

 

Repeat with methanol

Sample observations

 

Substance/alkanol Effect on adding equal amount of
ethanol/concentrated sulphuric(VI)acid
Absolute ethanol Sweet fruity smell
Methanol Sweet fruity smell

Explanation

 

Alkanols react with alkanoic acids to form a group of homologous series of sweet smelling compounds called esters and water. This reaction is catalyzed by concentrated sulphuric(VI)acid in the laboratory.

 

Alkanol           +    Alkanoic acid     Conc. H2SO4->     Ester     +  water

 

Naturally esterification is catalyzed by sunlight. Each ester has a characteristic smell derived from the many possible combinations of alkanols and alkanoic acids that create a variety of known natural(mostly in fruits) and synthetic(mostly in juices) esters .

 

Esters derive their names from the alkanol first then alkanoic acids. The alkanol becomesan alkyl group and the alkanoic acidbecomes” alkanoate hence alkylalkanoate. e.g.

 

Ethanol + Ethanoic acid -> Ethylethanoate + Water
Ethanol + Propanoic acid -> Ethylpropanoate + Water
Ethanol + Methanoic acid -> Ethylmethanoate + Water
Ethanol + butanoic acid -> Ethylbutanoate + Water
Propanol + Ethanoic acid -> Propylethanoate + Water
Methanol + Ethanoic acid -> Methyethanoate + Water
Methanol + Decanoic acid -> Methyldecanoate + Water
Decanol + Methanoic acid -> Decylmethanoate + Water
12

 

 

During the formation of the ester, the O joining the alkanol and alkanoic acid comes from the alkanol.

 

R1  -COOH  +      R2 OH

 

->

 

R1

 

-COO –R2

 

 

H2O

 

 

e.g.

 

  1. Ethanol reacts with ethanoic acid to form the ester ethyl ethanoate and water.

Ethanol            + Ethanoic acid         –Conc. H2SO4 –>Ethylethanoate + Water

C2H5OH (l)   + CH3COOH(l) –Conc. H2SO4 –>  CH3COO C2H5(aq)  +H2O(l)

CH3CH2OH (l)+ CH3COOH(l) –Conc. H2SO4 –>  CH3COOCH2CH3(aq)  +H2O(l)

 

  1. Ethanol reacts with propanoic acid to form the ester ethylpropanoate and water.

Ethanol            + Propanoic acid    –Conc. H2SO4 –>Ethylethanoate + Water

 

C2H5OH (l)+ CH3 CH2COOH(l) –Conc. H2SO4 –>CH3CH2COO C2H5(aq)  +H2O(l)

CH3CH2OH (l)+ CH3 CH2COOH(l) –Conc. H2SO4 –>

CH3 CH2COOCH2CH3(aq)  +H2O(l)

 

  1. Methanol reacts with ethanoic acid to form the ester methyl ethanoate and water. Methanol + Ethanoic acid –Conc. H2SO4 –>Methylethanoate + Water CH3OH (l) + CH3COOH(l) –Conc. H2SO4 –> CH3COO CH3(aq) +H2O(l)

 

 

  1. Methanol reacts with propanoic acid to form the ester methyl propanoate and water.

 

Methanol + propanoic acid –Conc. H2SO4 –>Methylpropanoate + Water CH3OH (l)+ CH3 CH2COOH(l) –Conc. H2SO4 –> CH3 CH2COO CH3(aq) +H2O(l)

 

  1. Propanol reacts with propanoic acid to form the ester propylpropanoate and water.

Propanol        + Propanoic acid      –Conc. H2SO4 –>Ethylethanoate + Water

C3H7OH (l)+ CH3 CH2COOH(l) –Conc. H2SO4 –>CH3CH2COO C3H7(aq)  +H2O(l)

CH3CH2 CH2OH (l)+ CH3 CH2COOH(l) –Conc. H2SO4 –>

CH3 CH2COOCH2 CH2CH3(aq)  +H2O(l)

 

(j)Oxidation

Experiment

 

Place 5cm3 of absolute ethanol in a test tube.Add three drops of acidified potassium manganate(VII).Shake thoroughly for one minute/warm.Test the solution mixture using pH paper. Repeat by adding acidified potassium dichromate(VII).

 

Sample observation table

Substance/alkanol Adding acidified pH of resulting Nature of resulting
KMnO4/K2Cr2O7 solution/mixture solution/mixture
Pure ethanol (i)Purple colour of pH= 4/5/6 Weakly acidic
13

 

 

KMnO4decolorized
(ii) Orange colour pH = 4/5/6 Weakly acidic
of K2Cr2O7turns
green.

Explanation

Both acidified KMnO4 and K2Cr2O7 are oxidizing agents(add oxygen to other compounds. They oxidize alkanols to a group of homologous series called alkanals then further oxidize them to alkanoic acids.The oxidizing agents are themselves reduced hence changing their colour:

 

  • Purple KMnO4 is reduced to colourless Mn2+ (ii)Orange K2Cr2O7is reduced to green Cr3+

The pH of alkanoic acids show they have few H+ because they are weak acids i.e

 

Alkanol           +              [O]    ->            Alkanal          +            [O]       ->          alkanoic acid

 

NB The [O] comes from the oxidizing agents acidified KMnO4 or K2Cr2O7 Examples

 

1.When ethanol is warmed with three drops of acidified KMnO4 there is decolorization of KMnO4

 

 

Ethanol + CH3CH2OH +

 

 

[O] -> Ethanal [O] -> CH3CH2O

 

 

+

+

 

 

[O]

 

[O]

 

 

->

 

->

 

 

Ethanoic acid

 

CH3COOH

 

 

 

2.When methanol is warmed with three drops of acidified K2Cr2O7 ,the orange

colour of acidified K2Cr2O7 changes to green.
methanol + [O] -> methanal + [O] -> methanoic acid
CH3OH + [O] -> CH3O + [O] -> HCOOH

 

3.When propanol is warmed with three drops of acidified K2Cr2O7 ,the orange colour

of acidified K2Cr2O7 changes to green.
Propanol   +    [O]   -> Propanal+   [O]    -> Propanoic acid
CH3CH2 CH2OH  +   [O] ->  CH3CH2 CH2O  +  [O] -> CH3 CH2COOH

4.When butanol is warmed with three drops of acidified K2Cr2O7 ,the orange colour of acidified K2Cr2O7 changes to green.

Butanol    +       [O]    ->       Butanal         +    [O]       ->   Butanoic acid

CH3CH2 CH2 CH2OH + [O]  ->CH3CH2 CH2CH2O +[O] -> CH3 CH2COOH

 

 

 

 

 

14

 

 

Air slowly oxidizes ethanol to dilute ethanoic acid commonly called vinegar. If beer is not tightly corked, a lot of carbon(IV)oxide escapes and there is slow oxidation of the beer making it flat.

 

(k)Hydrolysis /Hydration and Dehydration

 

  1. Hydrolysis/Hydration is the reaction of a compound/substance with water. Alkenes react with water vapour/steam at high temperatures and high pressures in presence of phosphoric acid catalyst to form alkanols.i.e.
Alkenes + Water – H3PO4 catalyst->  Alkanol
Examples

 

(i)Ethene is mixed with steam over a phosphoric acid catalyst at 300oC temperature and 60 atmosphere pressure to form ethanol

Ethene + water —60 atm/300oC/ H3PO4 –> Ethanol
H2C =CH2 (g)  +  H2O(l)  –60 atm/300oC/ H3PO4 –>  CH3 CH2OH(l)
This is the main method of producing large quantities of ethanol instead of
fermentation
(ii) Propene + water —60 atm/300oC/ H3PO4 –> Propanol
CH3C =CH2 (g)  +  H2O(l)  –60 atm/300oC/ H3PO4 –>  CH3 CH2 CH2OH(l)
(iii) Butene + water —60 atm/300oC/ H3PO4 –> Butanol

CH3 CH2 C=CH2 (g) + H2O(l)  –60 atm/300oC/ H3PO4  –> CH3 CH2 CH2 CH2OH(l)

 

  1. Dehydration is the process which concentrated sulphuric(VI)acid (dehydrating agent) removes water from a compound/substances.

 

Concentrated sulphuric(VI)acid dehydrates alkanols to the corresponding alkenes at about 180oC. i.e

 

Alkanol –Conc. H2 SO4/180oC–> Alkene + Water Examples

 

  1. At 180oC and in presence of Concentrated sulphuric(VI)acid, ethanol undergoes dehydration to form ethene.

Ethanol                    —180oC/ H2SO4   –>    Ethene                    +     Water

CH3 CH2OH(l)        –180oC/ H2SO4       –>    H2C =CH2 (g)  +  H2O(l)

  1. Propanol undergoes dehydration to form propene.

 

Propanol                    —180oC/ H2SO4    –>    Propene                    +                                         Water

CH3 CH2 CH2OH(l) –180oC/ H2SO4       –>    CH3CH =CH2 (g)  +  H2O(l)

  1. Butanol undergoes dehydration to form Butene.

Butanol                    —180oC/ H2SO4   –>    Butene                     +   Water

CH3 CH2 CH2CH2OH(l)         –180oC/ H2SO4       –>    CH3 CH2C =CH2 (g)

 

H2O(l)

  1. Pentanol undergoes dehydration to form Pentene.

 

Pentanol                    —180oC/ H2SO4    –>

 

 

 

 

Pentene

 

 

 

 

+   Water

 

 

 

15

 

CH3 CH2 CH2 CH2 CH2OH(l)–180oC/ H2SO4–>CH3 CH2 CH2C =CH2 (g)+H2O(l)

 

(l)Similarities of alkanols with Hydrocarbons I. Similarity with alkanes

 

Both alkanols and alkanes burn with a blue non-sooty flame to form carbon(IV)oxide(in excess air/oxygen)/carbon(II)oxide(in limited air) and water. This shows they are saturated with high C:H ratio. e.g.

 

Both ethanol and ethane ignite and burns in air with a blue non-sooty flame to form carbon(IV)oxide(in excess air/oxygen)/carbon(II)oxide(in limited air) and water.

 

CH2 CH2OH(l) + 3O2(g) -Excess air-> 2CO2 (g) +  3H2 O(l)
CH2 CH2OH(l) + 2O2(g) -Limited air->   2CO (g) +  3H2 O(l)
CH3 CH3(g) + 3O2(g) -Excess air-> 2CO2 (g) +  3H2 O(l)
2CH3 CH3(g) +  5O2(g) -Limited air-> 4CO (g)  +  6H2 O(l)

 

  1. Similarity with alkenes/alkynes

Both alkanols(R-OH) and alkenes/alkynes(with = C = C = double and C = C-triple ) bond:

 

(i)decolorize acidified KMnO4

(ii)turns Orange acidified K2Cr2O7  to green.

Alkanols(R-OH) are oxidized to alkanals(R-O) ant then alkanoic acids(R-OOH).

 

Alkenes are oxidized to alkanols with duo/double functional groups.

 

Examples

1.When ethanol is warmed with three drops of acidified K2Cr2O7 the orange of acidified K2Cr2O7 turns to green. Ethanol is oxidized to ethanol and then to ethanoic acid.

 

 

Ethanol + CH3CH2OH +

 

 

[O] -> Ethanal [O] -> CH3CH2O

 

 

+

 

+

 

 

[O]

 

[O]

 

 

-> Ethanoic acid -> CH3COOH

 

 

 

2.When ethene is bubbled in a test tube containing acidified K2Cr2O7 ,the orange of acidified K2Cr2O7 turns to green. Ethene is oxidized to ethan-1,2-diol.

Ethene            +    [O]            ->          Ethan-1,2-diol.

H2C=CH2        +    [O]            ->         HOCH2 -CH2OH

 

  • Differences with alkenes/alkynes

Alkanols do not decolorize bromine and chlorine water.

 

Alkenes decolorizes bromine and chlorine water to form halogenoalkanols Example

 

16

 

 

When ethene is bubbled in a test tube containing bromine water,the bromine water is decolorized. Ethene is oxidized to bromoethanol.

 

 

Ethene + Bromine water -> H2C=CH2 + HOBr ->

 

Bromoethanol.

 

BrCH2 -CH2OH

 

 

 

  1. Differences in melting and boiling point with Hydrocarbons Alkanos have higher melting point than the corresponding hydrocarbon (alkane/alkene/alkyne)

 

This is because most alkanols exist as dimer.A dimer is a molecule made up of two other molecules joined usually by van-der-waals forces/hydrogen bond or dative bonding.

 

Two alkanol molecules form a dimer joined by hydrogen bonding.

Example

 

In Ethanol the oxygen atom attracts/pulls the shared electrons in the covalent bond more to itself than Hydrogen.

 

This creates a partial negative charge (δ) on oxygen and partial positive charge(δ+) on hydrogen.

 

Two ethanol molecules attract each other at the partial charges through Hydrogen bonding forming a dimmer.

 

 

 

H H H Hydrogen bonds
H C C O H H Covalent bonds
H H H O C C H
H H

 

Dimerization of alkanols means more energy is needed to break/weaken the Hydrogen bonds before breaking/weakening the intermolecular forces joining the molecules of all organic compounds during boiling/melting.

 

E.USES OF SOME ALKANOLS

 

(a)Methanol is used as industrial alcohol and making methylated spirit (b)Ethanol is used:

 

  1. as alcohol in alcoholic drinks e.g Beer, wines and spirits. 2.as antiseptic to wash woulds

3.in manufacture of vanishes, ink ,glue and paint because it is volatile and thus

easily evaporate

4.as a fuel when blended with petrol to make gasohol.

 

17

 

 

 

 

 

 

 

B.ALKANOIC ACIDS (Carboxylic acids)

 

(A) INTRODUCTION.

 

Alkanoic acids belong to a homologous series of organic compounds with a general formula CnH2n +1 COOH and thus -COOH as the functional group .The 1st ten alkanoic acids include:

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

18

 

 

Alkanoic acids like alkanols /alkanes/alkenes/alkynes form a homologous series

where:
(i)the general name of an alkanoic acids is derived from the alkane name then
ending with “–oic” acid as the table above shows.
(ii) the members have R-COOH/R C-O-H as the functional  group.
n General /molecular Structural formula IUPAC
formular name
0 HCOOH HCO-H Methanoic
acid
O
1 CH3 COOH H Ethanoic
acid
HC CO-H
O
H
2 CH3 CH2 COOH H H Propanoic
C2 H5 COOH acid
H-CC COH
H O
H
3 CH3 CH2 CH2 COOH H H H Butanoic
C3 H7 COOH acid
H-C- CC COH
H O
H H
4 CH3CH2CH2CH2 H H H H Pentanoic
COOH acid
C4 H9 COOH H- CC- CC COH
H H O
H H
5 CH3CH2 CH2CH2CH2 H H H H H Hexanoic
COOH acid
C5 H11 COOH H C- C C- C C COH
H H O
H H H
6 CH3CH2 CH2 H H H H H H Pentanoic
CH2CH2CH2 COOH acid
C6 H13 COOH H  C C- CC- CCCOH
19
H H H H O