Home Teachers' Resources CHEMISTRY FORM 1, 2, 3 & 4 COMPREHENSIVE NOTES

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CHEMISTRY FORM 1, 2, 3 & 4 COMPREHENSIVE NOTES

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November 25, 2024

UNIT 1: ATOMIC STRUCTURE AND THE PERIODIC TABLE.

Table of Contents

Checklist

Structure of the atom
The subatomic particles;

Protons
Neutrons
Electrons;

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CHEMISTRY FORM ONE NOTES

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Atomic number and mass number
Isotopes;
Energy levels and electron arrangements.
The periodic table;

Groups
Periods;

Relative atomic mass and isotopes;
Ion formation;

Ion;
Cations;
Anions;
Ionization energy;
Electron affinity;

Valency and oxidation numbers;
Chemical formula;
Chemical equations
Balancing chemical equations;

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The atom:

– Refers to the smallest particle of an element hat can take part in a chemical reaction;

– It has an average diameter of 10^-8 cm with a nucleus of about 10^-13 cm;

Parts of an atom

– The atom is made of two main parts:

The nucleus
The energy levels;

The nucleus:

– Is the positively charged part of an atom;

– The nucleus contains two subatomic particles; neutrons and protons;

– The positive charge is due to presence of protons;

– The nuclei of all atoms contain neutrons except the hydrogen atom;

– The protons and the neutrons are together referred to as the nucleons;

The energy levels.

– They contain the electrons;

– Electrons are so small and move so fast that their path cannot be traced directly;

– Thus the energy level simple represents the region where the electrons are most likely to be found;

Structure of the atom.

Note:

– The atom can still however be split into smaller particles termed the sub-atomic particles;

The sub-atomic particles.

– Are generally three:

Protons;
Neutrons;
Protons;

Protons.

– Are the positively charged sub-atomic particles;

– Are found in the nucleus and thus form part of the nucleons;

– The number of protons in the nucleus is equal to the number of electrons in the energy levels;

Neutrons.

– Are neutrally charged sub-atomic particles found in the nucleus of the atom;

– They are thought to probably prevent the positively charged protons from getting too close to each other;

Electrons.

– Are negatively charged sub-atomic particles found in the energy levels;

– The number of electrons in the energy levels is equal to the number of protons in the nucleus;

– This makes the atom to be electrically neutral;

Atomic number and mass number.

Atomic number.

– Refers to the number of protons in the nucleus of an atom;

Examples.

Sodium has 11 protons I the nucleus and thus said to have atomic number 11;
Chlorine has 17 protons in the nucleus and thus said to have atomic number 17;

Mass number;

– Refers to the sum of the number of protons and neutrons in an atom of an element;

Examples:

Sodium has 2 neutrons and 11 protons hence a mass number of 23;
Chlorine has 18 neutrons and 17 protons hence a mass number of 35.

Notation of atomic number and mass number;

– Both atomic number and mass number of an element can be written along with the symbol of an element;

Mass number;

– Is conventionally represented as a superscript to the left of the symbol;

Examples:

Sodium; ²³Na;

Magnesium ²⁴Mg;

Atomic number;

– Is conventionally represented as a superscript to the left of the symbol;

Examples:

Sodium; ₁₁Na;

Magnesium ₁₂Mg;

Thus the elements can be conventionally represented as:

Sodium ²³Na

Magnesium ²⁴Mg

Atomic properties of the first 20 elements.

Element	Symbol	Number of electrons	Number of protons	Number Of neutrons	Atomic number	Mass number
Hydrogen	H
Helium	He
Lithium	Li
Beryllium	Be
Boron	B
Carbon	C
Nitrogen	N
Oxygen	O
Fluorine	F
Neon	Ne
Sodium	Na
Magnesium	Mg
Aluminium	Al
Silicon	Si
Phosphorus	P
Sulphur	S
Chlorine	Cl
Argon	Ar
Potassium	K
Calcium	Ca

Isotopes.

– Are atoms of the same element with same atomic number but different mass number due to different number of neutrons.

Examples of isotopes.

Element	Isotope	Atomic No.	Number of protons	Number of neutrons	Mass number	Isotopic representation
Hydrogen


Carbon
Carbon
Oxygen


Chlorine
Chlorine

Energy levels and electron arrangements.

Energy levels:

– Are definite orbits in an atom that the electrons occupy.

– The energy levels are numbered 1, 2, 3starting with the one closest to the nucleus.

– Electrons occupying the same energy level have approximately the same amount of energy.

– Each energy level can only accommodate a given maximum number of electrons.

Maximum number of electrons per energy level

Energy level	Maximum number of electrons
1^st	2
2^nd	8
3^rd	8 (only for the first 20 elements)

Illustrations:

Hydrogen

– It has only one electron and thus this electron occupies the first energy level.

– Since the first energy level is not yet full, hydrogen does not have the second energy level;

– The electron arrangement of hydrogen is thus 1.

Helium:

– Helium is atomic number 2 and has only two electrons, which occupy the first energy level.

– The first energy level is thus completely full, but since there are no other r electrons lithium also has only one energy level;

– The electron arrangement is thus 2.

Chlorine:

– Chlorine has atomic number 17 and thus has 17 electrons;

– The first two electrons occupy the fist energy level which is thus completely filled up;

– The remaining 15 electrons occupy the second energy level, which can however accommodate only 8 to be completely filled up;

– Thus the remaining 7 electrons move to the third energy level; which needs 8 to be completely filled up;

– Since the third energy level is not yet full chlorine does not have a fourth energy level;

– The electron arrangement is thus 2.8.7.

Electron arrangement.

– Refers to the distribution of electrons in the energy levels of an atom.

Example: electron arrangement for the first 20 elements.

Element	Symbol	Atomic number	No. of electrons	Electron arrangement
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulphur
Chlorine
Argon
Potassium
Calcium

Dot and cross diagrams.

– Is a diagrammatic representation of the electron arrangements in an atom in which the energy levels are represented by concentric lines while electrons are represented by dots or crosses.

– However all electrons are the same regardless of whether they are represented as dots or crosses.

Examples:

Lithium	Magnesium	Aluminium	Carbon

No. of electrons =	No. of electrons	No. of electrons	No. of electrons
No. of protons =	No. of protons	No. of protons	No. of protons
No. of neutrons =	No. of neutrons	No. of neutrons	No. of neutrons

Beryllium	Nitrogen	Chlorine	Argon

No. of electrons =	No. of electrons =	No. of electrons =	No. of electrons =
No. of protons =	No. of protons =	No. of protons =	No. of protons =
No. of neutrons =	No. of neutrons =	No. of neutrons =	No. of neutrons =

Potassium	Boron	Silicon	Phosphorus

No. of electrons =	No. of electrons =	No. of electrons =	No. of electrons =
No. of protons =	No. of protons =	No. of protons =	No. of protons =
No. of neutrons =	No. of neutrons =	No. of neutrons =	No. of neutrons =

The Periodic table

– Is a table showing the arrangements of elements in order of their relative atomic masses.

– Is based on the ideas of Dmitri Ivanovich Medeleev.

– The modern periodic table is based on Meneleevs periodic law which states that:

The properties of elements are a periodic functions of their relative atomic masses

– The modern periodic law itself states that:

The properties of elements are a periodic functions of their atomic numbers

Design of the modern periodic table.

It has vertical columns called groups and horizontal rows called periods.

Groups:

– Are the vertical columns of a periodic table.

– Are eight in number; ad numbered in capital Roman numerals I all through to VIII.

Note: group VIII is also called group zero because the elements have little tendency to gain o lose electrons during chemical reactions

– Between group 2 and group 3 is a group of elements called the transition metals;

Periods:

– Are the horizontal rows in a periodic table.

– They are 8 in number in a modern periodic table.

Transition metals.

– Are elements that form a shallow rectangle between group II and group III.

– These elements are generally metallic and hence the name transition metals”

– They are not fitted in any group because they have variable valencies.

– Are sometimes called the d-block elements.

– They are much less reactive than the elements in groups I and II

– They have some unique characteristics that make them not fit in the 8 groups of the periodic table

– These include:

Have variable valencies; hence show different oxidation states in their compounds;
They form coloured compounds as solids and in aqueous solutions;
Have very high melting and boiling points (than metals in groups I and II).
They do not react with water;
Have very high densities (compared to metals in groups I ad II)

Lanthanides and Actinides.

– They form a block of elements within the transition metals.

– Are sometimes called the inner transition metals.

– The lanthanides consist of 14 elements from Cerium (Ce) to Lutetium (Lu).

– The Actinides are the 14 elements from Thorium (Tho) to Lawrencium (Lr).

Placing an elements in the periodic table

– The position of an elements in the periodic table is governed by the atomic number and hence the

electron arrangement.

The period:

– The period to which an element belongs is determined by the number of energy levels.

– The number of energy levels is equal to the period to which an elements belongs.

Examples:

Elements	Symbol	Atomic number	Electron arrangement	Number of energy levels	Period
Lithium	Li	3	2.1	2	2
Sodium	Na	11	2.8.1	3	3
Calcium	Ca	20	2.8.8.2	4	4
Nitrogen	N	7	2.5	2	2
Helium	He	2	2	1	1

Group:

– The group to which an element belongs to is governed by the number of electrons in the

outermost energy level.

– The number of electrons in the outermost energy level is equal to the group to which the element

belongs.

Examples:

Elements	Symbol	Atomic number	Electron arrangement	Outermost electrons	Group
Potassium	K	19	2.8.8.1	1	1
Aluminium	Al	13	2.8.3	3	3
Silicon	Si	14	2.8.4	4	4
Oxygen	O	8	2.6	6	6
Chlorine	Cl	17	2.8.7	7	7

Diagram of a full periodic table.

Note: In the modern periodic table atomic masses are used instead of mass numbers. The atomic

masses are preferable because they take care of elements with isotopes unlike mass numbers.

Diagram: part of the periodic table showing the first 20 elements

Relative Atomic Mass and Isotopes.

Introduction:

– The masses of individual atoms of elements are very negligible and thus quite difficult to weigh.

– On average the mass of an atom is approximately 10^-22g which cannot be determined by an

ordinary laboratory balance.

– For this reason the mass of atom has been expressed relative to that of a chosen standard element

hence the term relative atomic mass.

– The initial reference element was hydrogen which was later replaced with oxygen.

– Later the oxygen scale was found unsuitable;

Reason:

Oxygen exists I several isotopes and thus led to problems when deciding the mass of an oxygen atom.

– For this reason oxygen was replaced with carbon as the reference atom and to date relative atomic

masses of elements are based on an atom of carbon-12 (note that carbon is isotopic and exists as

Carbon -12 or carbon-14).

Definition:

Relative atomic mass (R.A.M) of an element refers to the average mass of an atom of the element

compared with a twelfth (1/12) of an atom of carbon-12.

RAM = Average mass of one atom of an element

¹/₁₂ mass of an atom of carbon-12

Measurement of Relative atomic mass

– RAM of elements is determined by an instrument called Mass Spectrometer.

– The instrument can also be used to determine the relative abundance of isotopes.

– The use of a mass spectrometer in determining the RAM of elements is called mass spectrometry.

How mass spectrometry works.

– In the mass spectrometer atoms and molecules are converted into ions.

– The ions are then separated as a result of the deflection which occurs in a magnetic field.

– Each ion (from an atom, isotope or molecule) gives a deflection which is amplified into a trace.

– The height of each peak measures the relative abundance of the ion which gives rise to that peak.

Note:

– Generally the relative atomic mass of an element is closest in value to the mass of the most

abundant isotope of the element.

Example: Diagram of a spectrometer trace for Lithium

Explanations:

– The trace has two peaks indicating that there are two isotopes for lithium.

– The fist peak occurs at a relative isotopic mass of 6 and the second at 7; these are the RAM of the two

isotopes respectively.

– The percentage abundance of the isotope with RAM of 6 (⁶Li) is 9 while the RAM of the isotope with

RAM 7 (⁷Li) is 91.

Calculating relative atomic masses of isotopic elements.

– Information form a spectrometer trace is usually extracted and used in calculation the relative atomic

mass of elements.

Worked examples.

The mass spectrum below shows the isotopes present in a sample of lithium.

(i). Use this mass spectrum to help you complete the table below for each lithium isotope in the sample. (3 marks)

Isotope	Percentage composition	Number of
Isotope	Percentage composition	Protons	Neutrons
⁶Li
⁷Li

(ii). Calculate the relative atomic mass of this lithium sample. Your answer should be given to three significant figures. (3 marks)

Element X with atomic number 16 has two isotopes. ⅔ of ³³X and ⅓ of ³⁰X. What is the relative atomic mass of element X? (2 marks)
Calculate the relative atomic mass of an element whose isotopic masses and relative abundances are shown below. (2 marks)

A neutral atom of silicon contains 14 electrons, 92% of silicon – 28, 5% silicon 29 and 3% silicon 30

(i). What is the atomic number of silicon? (1mark)

(ii). Calculate the relative atomic mass of silicon. (1mark)

Oxygen exists naturally as isotopes of mass numbers 16, 17 and 18 in the ratio 96:2:2 respectively. Calculate its R.A.M (2 marks)

Calculate the relative atomic mass of potassium from the isotopic composition given below.

Isotope Relative abundance

³⁹K 93.1

⁴⁰K 0.01

⁴¹K 6.89

Sulphur and sulphur compounds are common in the environment.

(a). A sample of sulphur form a volcano contained 88% by mass of ³²S and 12% by mass of ³⁴S.

(i). Complete the table below to show the atomic structure of each isotope of sulphur.

Isotope	Number of
Isotope	Protons	Neutrons	Electrons
³²S
³⁴S

(ii). Define relative atomic mass. (2 marks)

(iii). Calculate the relative atomic mass of the volcanic sulphur. (2 marks)

Iridium, atomic number 77, is a very dense metal. Scientists believe that meteorites have deposited virtually all the iridium present on earth. A fragment of a meteorite was analysed using a mass spectrometer and a section of the mass spectrum showing the isotopes present in iridium is shown below.

(a). Explain the term isotopes. (1mark)

(b). Use the mass spectrum to help you complete the table below for each iridium isotope in the meteorite.

Isotope	Percentage composition	Number of
Isotope	Percentage composition	Protons	Neutrons
⁶Ir
⁷Ir

(ii). Calculate the relative atomic mass of the iridium in this meteorite. (3 marks)

Ion formation.

Introduction:

– Atoms whose outermost energy levels contain the maximum possible number of electrons are said to be stable.

– Thus atoms with energy levels 2, 2.8 and 2.8.8 are said to be stable.

– Electron configuration 2 is said to have a stable duplet state while electron configuration 2.8 and 2.8.8 is said to have a stable octet state.

– These electron configurations resemble those of noble gases and as such they are stable and do not react.

– Atoms without this stability acquire it by either electron gain or electron loss.

– Whether an atom loses or gains electro(s) depend on the number of electrons in the outermost energy level.

– Take the case of sodium.

– Atomic number is 11 with an electron configuration of 2.8.1.

– Thus sodium has two options in to become stable:

to lose the single electron and acquire a stable electron configuration of 2.8.
to gain 7 electrons in its outermost energy level and acquire a stable electron configuration of 2.8.8

– Gaining a single electrons and losing a single electrons requires equal amounts of energy.

– Thus it is cheaper and faster in terms of energy for sodium to lose the single electron in the outermost energy level than to gain 7 electrons into its outermost energy level.

– Thus sodium acquires a stable electron configuration 2.8 by losing the single electron in its outermost energy level.

Diagram

Sodium atom Sodium ion.

Equation:

Na → Na⁺ + e^–

Further examples:

Element	Electron arrangement	Options for stability	Best (cheapest) option
Chlorine	2.8.7	2.8 or 2.8.8	2.8.8
Potassium	2.8.8.1	2.8.8 or 2.8.8.8	2.8.8
Aluminium	2.8.3	2.8 or 2.8.8	2.8
Magnesium	2.8.2	2.8 or 2.8.8	2.8
Carbon	2.4	2 or 2.8	2. or 2.8
Oxygen	2.6	2. or 2.8	2.8

Ion:

Definition: an ion is a charged particle of an element.

– Are formed when an atom of an element either loses or gains electrons.

Illustration:

– For a neutral atom the number of electrons in the energy levels (negative charges) is equal and thus

completely balances the number of protons in the nucleus (positive charges).

– Thus the net charge in a neutral atom is zero (0).

– When an atom gains electron(s), the number of electrons becomes higher than the number of protons

resulting to a net negative charge hence an ion.

– Oppositely when an atom loses electron(s) the number of protons becomes higher than the number of

electrons resulting into a net positive charge hence an ion.

– The charge on the ion is usually indicated as a superscript to the right of the chemical symbol.

– Thus ions are of two types:

Cations
Anions

Cations:

– Are positively charged ions.

– Are formed when atoms lose electrons resulting into the number of protons being higher than the number of electrons.

– Are mostly ions of metallic elements since most metals react by electron loss.

Examples:

(i). Magnesium:

– It has atomic number 12, with electron arrangement 2.8.2.

– It has 12 protons and 12 electrons hence a net charge of 0 hence the atom is written simply as Mg.

– It will form its ions by losing the two electrons from the outermost energy level.

– Thus the number of electrons decreases to 10 while the number of protons remains 12.

– This leads to a net charge of +2, giving the ion with the formula Mg²⁺.

Diagrammatic illustration:

(ii). Phosphorus

– It has atomic number 13, with electron arrangement 2.8.3.

– It has 13 protons and 13 electrons hence a net charge of 0 hence the atom is written simply as Al.

– It will form its ions by losing three (3) electrons out of the outermost energy level.

– Thus the number of electrons decreases by three to 10 while the number of protons remains 13.

– This leads to a net charge of +3, giving the ion with the formula Al⁺³.

Diagrammatic illustration:

Anions:

– Are negatively charged ions.

– Are formed when atoms gain electrons resulting into the number of electrons being higher than the number of protons.

– Are mostly ions of non-metallic elements since most non-metals ionize (react) by electron gain.

Examples.

(i). Chlorine:

– It has atomic number 17, with electron arrangement 2.8.7.

– It has 17 protons and 17 electrons hence a net charge of 0 hence the atom is written simply as Cl.

– It will form its ions by gaining a single electron into the outermost energy level.

– Thus the number of electrons increases to 18 while the number of protons remains 17.

– This leads to a net charge of -1, giving the ion with the formula Cl^–.

Diagrammatic illustration:

Note: the electros gained must be represented by a different notation from the initial electrons in the atom. E.g. if the initial electrons are represented with crosses (x) then the gained electrons should be represented by dots (.) and vise versa.

(ii). Phosphorus

– It has atomic number 15, with electron arrangement 2.8.5.

– It has 15 protons and 15 electrons hence a net charge of 0 hence the atom is written simply as P.

– It will form its ions by gaining three (3) electrons into the outermost energy level.

– Thus the umber of electrons increases by three to 18 while the number of protons remains 15.

– This leads to a net charge of -2, giving the ion with the formula P^-3.

Diagrammatic illustration:

Electron transfer during chemical reactions.

– Atoms react either by electron gain or electron loss.

– Generally metals react by electron gain while non-metals react by electron loss.

Illustration:

– Consider the reaction between sodium and chlorine.

– Sodium attains stability // reacts by losing the single electron form its outermost energy level.

– Chlorine attains stability // reacts by gaining a single electron into its outermost energy level.

– Thus during the reaction between the two elements the single electron lost by the sodium atom to

form the sodium ion is the same one gained by the chlorine atom to form the chloride ion.

Some basic concepts.

Valence electrons:

– Refers to the number electrons in the outermost energy level.

Examples:

– Calcium, with electron arrangement 2.8.8.2 has 2 valence electrons.

– Oxygen with electron arrangement 2.6 has 6 valence electrons.

– Phosphorus with electron arrangement 2.8.5 has 5 valence electrons.

Valency:

– Refers to the number of electrons an atom loses or gains during a chemical reaction.

– Valency is also known as the combining power of an element.

Examples:

– Calcium, with electron arrangement 2.8.8.2 loses 2 electrons during chemical reactions and hence has

a valency of 2.

– Oxygen with electron arrangement 2.6 gains 2 electrons during chemical reactions and thus has a

valency of 2.

– Phosphorus with electron arrangement 2.8.5 has gains 3 electrons during chemical reactions and

hence has a valency of 3.

– Aluminium, with electron arrangement 2.8.3 loses 3 electrons during chemical reactions and hence

has a valency of 2.

Note: Some elements have variable valencies and are usually termed the transitional elements (metals)

Examples:

– Iron can have valency 2 or 3;

– Copper can have valency 1 or 2

– Lead can have valency 2 or 4.

Summary on valencies of common elements.

Valency 1

Valency 2

Valency 3

Metals

Sodium

Potassium

Calcium

Barium

Magnesium

Zinc

Iron

Lead

Copper

Aluminium

Iron

Non-metals

Nitrogen

Chlorine

Fluorine

Hydrogen

Nitrogen

Oxygen

Sulphur

Nitrogen

Phosphorus

Radicals.

– Are groups of atoms with a net charge that exist and react as a unit during chemical reactions.

– Radicals also have a valency, which is equivalent to the value of its charge.

Summary on valencies of some common radicals.

Valency 1

Valency 2

Valency 3

Radicals

Ammonium (NH₄⁺)

Hydroxide (OH^–)

Nitrate (NO₃^–)

Hydrogen carbonate (HCO₃^–)

Hydrogen sulphate (HSO₄^–)

Carbonate (CO₃^2-)

Sulphate (SO₄^2-)

Sulphite (SO₃^2-)

Phosphate (PO₄^3-)

Oxidation number.

– Refers to the number of electrons an atom loses or gains during a chemical reaction.

– In writing the oxidation number the sign (+ or -) to show gain or loss is written followed by the

number of electrons lost or gained respectively.

Illustration.

– Atoms are electrically neutral and are thus assigned an oxidation state of 0 since the number of

protons in the nucleus is equal to the number of electrons in the energy levels.

– However when atoms react they either lose or gain electrons and thus acquire a new state.

– This new state is a new oxidation state and the atom thus acquires a new oxidation number

Examples:

Atom	E. arrangement	Ion formula	Valency	Oxidation number
Sodium	2.8.1	Na⁺	1	+1
Magnesium	2.8.2	Mg²⁺	2	+2
Aluminium	2.8.3	Al³⁺	3	+3
Nitrogen	2.5	N^3-	3	-3
Sulphur	2.8.6	S^2-	2	-2
Chlorine	2.8.7	Cl^–	1	-1

Further examples:

Particle	Oxidation number
Copper metal, Cu	0
Lead (II) ion, Pb²⁺	+2
Bromide ion, Br^–	-1
Aluminium ion, Al²⁺	+2
Sulphide ion, S^2-	-2
Magnesium metal, Mg	0

Note: oxidation number (state) and charge of an element.

– Oxidation state is written with the positive or the negative sign coming before the element.

Examples: -2, 3, +1, -1 etc.

– Charge on an element is write as a superscript of the element with the number coming before the

positive r the negative sign

Examples: Mg²⁺, Al³⁺, Na⁺, Cl^– etc.

Chemical formulae.

– Refers to a representation of a chemical substance using chemical symbols.

– In a single atom it is equivalent to the chemical symbol of the element.

– In a compound it shows the constituent elements and the proportions in which they are combined.

Deriving the chemical formula of compounds.

– In order to write the correct formula of a compound the following must be known:

The symbols of the constituent elements or radicals.
The valencies of the elements or radicals

– The chemical formula should start with the element which is more likely to lose electron (s) followed

by the element that is more likely to gain.

Worked examples.

Deriving the formula of sodium chloride.

Elements	Sodium	Chlorine
Formula	Na	Cl
valencies	1	1
balancing	x1	x1

Balancing ratios as subscripts: Na₁Cl₁

Formula: NaCl

Explanation:

– For sodium to combine with chlorine to form sodium chloride, sodium loses an electron while

chlorine gains an electron.

– Thus every sodium atom needs only a single chlorine atom for both t be fully stable

Note:

– When the balancing ratio // subscript is 1, it is usually not written since the symbol of the

element itself represents a single atom.

Deriving the formula of magnesium chloride.

Elements	Magnesium	Chlorine
Formula	Mg	Cl
valencies	2	1
balancing	X1	X2

Balancing ratios as subscripts: Mg₁Cl₂

Formula: MgCl₂

Deriving the formula of magnesium oxide.

Elements	Magnesium	Oxygen
Formula	Mg	O
valencies	2	2
balancing	X1	X1

Balancing ratios as subscripts: Mg₁O₁

Formula: MgO

Further worked examples.

Derive the chemical formula of each of the following compounds.

Calcium fluoride
Carbon (II) oxide
Carbon (IV) oxide
Aluminium nitrate
Calcium hydrogen carbonate

Complete the table below for elements A, B and C

Element	Valency	Chemical formula of various compounds
Element	Valency	hydroxides	Sulphates	carbonates	nitrates	phosphates	Hydrogen carbonates
A	1
B	2
B	3

Chemical equations.

– Refers to representations of a chemical reaction by means of chemical symbols and formula.

Key features of a chemical equation.

– The correct formulae of the reactants are on the left of the equation.

– The correct formulae of the products are on the right of the equation.

– The reactants and products are separated by an arrow pointing to the right.

– The state symbols of the reactants and products must be stated as subscripts to the right of the symbols

– The number of each atom on the reactants side must be equal to the number of the same atom on the products side.

Example.

Reaction between hot copper metals and oxygen gas.

Word equation: Copper + oxygen → Copper (II) oxide.

Chemical equation: 2Cu_(s) + O_2(g) → 2CuO_(s).

Balancing chemical equations.

– A chemical equation is only valid if it is balanced.

– A chemical equation is said to be balanced if the number of each atom on the reactants side is equal to that on the products side.

– This is because atoms are neither created nor destroyed during a chemical reaction.

Rules and guidelines in balancing chemical equations.

Step 1: Write the chemical equation in words.

Example: Copper metal + oxygen gas.

Step II: Write the correct formulae of both reactants and products

Example: Cu + O₂ → CuO

Step III: Check whether the number of atoms of each element on the reactants side is equal to that on the products side.

– If equal proceed to step (V);

– If not equal proceed to step (IV).

Example: Cu + O₂ → CuO

– In this case there are two oxygen atoms on the reactants side yet there is only one oxygen atom on the products side. Thus we proceed to step IV

Step IV: Multiply the chemical formula containing the unbalanced atoms with the lowest common multiple.

Example: Cu + O₂ → CuO

– In this case the chemical formula with the unbalanced atom is CuO on the products side. We thus multiply it by 2.

– The new equation now reads: Cu + O₂ → 2CuO

Step V: check again to ensure that all atoms are balanced.

– If all atoms are balanced proceed to step VI.

– If not then repeat step IV until all atoms are balanced.

Example: Cu + O₂ → 2CuO

– In this case multiplying CuO by 2 offsets the balancing of Cu; which is now unbalanced!

– We therefore repeat step IV in order to balance Cu.

– There is only 1 Cu atom on the reactants side yet there are 2 Cu atoms on the products side.

– We thus multiply the formula with the unbalanced atom (s) by the lowest common multiple, in this case 2.

– The new equation at this step thus becomes: 2Cu + O₂ → 2CuO
– We then repeat step V; in this case all atoms are now balanced.

Step VI: The physical states of the reactants and the products are then indicated.

– If this is not done the chemical equation is considered incorrect.

Types of state symbols.

– There four main state symbols.

Solid; denoted as (s)
Liquid; denoted as (l)
Aqueous (in solution in water); denoted as (aq)
Gaseous; denoted as (g)

– In a chemical equation the state symbols are written with their denotations as subscripts to the right of the chemical formulae.

Example: 2Cu_(s) + O_2(g) → 2CuO_(s)

Thus the balanced chemical equation for the reaction between copper metal ad oxygen is:

2Cu_(s) + O_2(g) → 2CuO_(s)

Worked examples:

Balance equations for each of the following reactions.

Sodium hydroxide and dilute hydrochloric acid

Zinc oxide and dilute sulphuric (VI) acid

Zinc metal and dilute nitric (V) acid

Calcium carbonate and dilute sulphuric (VI) acid

Sodium and water

Balance each of the following equations.

Mg_(s) + HCl_(aq) →MgCl_2(aq) + H_2(g)

Na_(s) + H₂O_(l) → NaOH_(aq) + H_2(g)

NaOH_(aq) + H₂SO_4(aq) →Na₂SO_4(aq) + H₂O_(l)

CuCO_3(s) + HNO_3(aq) → Cu(NO₃)_2(aq) + CO_2(g) + H₂O_(l)

H₂S_(g) + O_2(g) → O_2(g) + H₂O_(l)

C₂H_6(g) + O_2(g) → CO_2(g) + H₂O_(l)

Pb(NO₃)_2(s) → PbO_(s) + NO_2(g) + O_2(g)

Fe_(s) + Cl_2(g) → FeCl_3(s)

Al_(s) + H₂SO_4(aq) → Al₂(SO₄)_3(aq) + H_2(g).

UNIT 2: CHEMICAL FAMILIES; PATTERNS AND PROPERTIES.

Checklist.

– Meaning of chemical families.

– Main chemical families;

Alkali metals
- Meaning and members
- Trends down the group
  - Atomic radius
  - Ionic radius
  - Energy levels
- Physical properties
  - Appearance
  - Ease of cutting
  - Melting and boiling [points
  - Electrical conductivity
  - 1^st ionization energy
- Chemical properties
  - Burning in air
  - Exposure to air
  - Reaction with water
  - Reaction with chlorine
- Similarity of ions and formula of compounds of alkali metals
- Uses of alkali metals.
Alkaline earth metals
- Meaning and members
- Trends down the group
  - Atomic radius
  - Ionic radius
  - Energy levels
- Physical properties
  - Melting and boiling [points
  - Electrical conductivity
  - 1^st ionization energy
  - 2^nd ionization energy
- Chemical properties
  - Reaction with air and water
  - Reaction with steam
  - Reaction with chlorine
  - Reaction with dilute acids
- Similarity of ions and formula of compounds of alkaline earth metals
- Uses of alkaline earth metals.
The Halogens
- Meaning and members
- Trends down the group
  - Atomic radius
  - Ionic radius
  - Energy levels
- Physical properties
  - Preparation and properties of chlorine
  - Appearance and physical states of halogens at room temperature
  - Melting and boiling points
  - Electrical conductivity
- Chemical properties
  - Ion formation
  - Electron affinity
  - Reaction with metals
  - Reaction with water
  - Reactivity trend of halogens.
- Similarity of ions and formula of compounds of alkali metals
- Uses of alkali metals.
The noble gases
- Meaning and members
- Trends down the group
  - Atomic radius
  - Ionic radius
  - Energy levels
- Physical properties
  - Melting and boiling [points
  - Electrical conductivity
  - 1^st ionization energy
- Chemical properties
- Uses of the noble gases

Introduction:

– Elements are classified and hence positioned in the periodic table based on the number of valence electrons and the number of energy levels.

– The number of valence electrons is equal to the group to which the element belongs; while the number of energy levels is equal to the period to which the element belongs.

– Elements in the same group are said to belong to the same chemical family.

– Thus a chemical family refers to a group of elements in the same number of valence electrons ad hence in the same group of the periodic table.

Characteristics of a chemical family:

– have same number of valence electrons;

– Show a uniform gradation in physical properties;

– have similar chemical properties;

Main chemical families.

– Four main chemical families will be studies in this section.

The Alkali metals
The Alkaline earth metals
The halogens
The noble gases

The Alkali metals.

– Are the elements with one valence electron and hence in group I of the periodic table.

– All are metallic in nature.

– The members of the family in order down the group is as follows:

Lithium
Sodium
Potassium
Rubidium
Caesium
Francium

Electron arrangement of the first three alkali metals.

Elements	Electron arrangement
Lithium	2.1
Sodium	2.8.1
Potassium	2.8.81

Diagram: Part of periodic table showing the alkali metals

Gradation in properties of alkali metals.

Atomic and ionic radius.

Atomic radius:

– Refers to the distance between the centre of the nucleus of an atom and the outermost energy level occupied by electron(s)

Ionic radius:

– Refers to the distance between the centre of the nucleus of an ion and the outermost energy level occupied by electron(s)

Trend:

– The ionic radius and the atomic radius of alkali metals increase down the group.

Reason:

– There is an increase in the number of energy levels down the group from lithium to Francium.

Illustration:

– Lithium (2.1) has only 2 energy levels; sodium (2.8.1) has 3 energy levels while potassium (2.8.8.1) has 4 energy levels.

– Thus the outermost electron in potassium is further from the nucleus than the outermost electron in sodium and lithium.

Atomic and ionic radius of the same element.

– For the same alkali metals the atomic radius is larger than the ionic radius.

Reason:

– Alkali metals form ions by losing the valence electron, lading to the loss of an entire outermost energy level. Thus the atoms have more energy levels than the corresponding ion hence a larger radius in the atom than I the ion.

Illustration

– Potassium atom has electron arrangement of 2.8.8.1 hence 4 energy levels.

– During ion formation potassium reacts by losing the single valence electron to acquire a new electron arrangement of 2.8.8 hence 3 energy levels.

– Thus the ion has a smaller radius than the atom.

Diagrammatically: Potassium atom and potassium ion

Summary: changes in atomic and ionic radius among alkali metals.

Element	Symbol	Atomic number	Atomic radius (nm)	Ionic radius (nm)
Lithium	Li	3	0.133	0.060
Sodium	Na	11	0.157	0.095
Potassium	K	19	0.203	0.133

Physical properties of alkali metals.

Appearance.

– Alkali metals have metallic luster when freshly cut. This refers to a shiny appearance on th cut surface.

– This surface however tarnishes due to reaction on exposure to air.

Ease of cutting.

– They are soft and easy to cut.

– The softness and ease of cutting increase down the group.

Reason:

– Alkali metals have giant metallic structures held together by metallic bonds.

– Metallic bond is due to attraction between the positively charged nucleus of one atom and the electrons in the outermost energy level of the next atom.

– Thus the force of attraction is stronger is smaller atoms than in larger atoms.

– The increase in atomic radius down the group implies that the strength of metallic bonds also decrease down the group (hence ease of cutting and softness).

They have relatively low melting and boiling points (in comparison to other metals).

Reason: they have relatively weaker metallic bonds.

– The melting and boiling points decrease down the group.

Reason: – The size of the atoms increase down the group due to increasing number of energy levels hence decrease in the strength of the metallic bonds (down the group).

Electrical conductivity.

– Alkali metals are good conductors of heat ad electricity.

Reason: they have delocalized electrons in the outermost energy level.

– The electrical conductivity is similar for all alkali metals.

Reason: all alkali metals have the same number of delocalized electron (a single electron) in the outermost energy level.

Note:

– In metals the electrons in the outermost energy level do not remain in one fixed position. They move randomly throughout the metallic structure and are thus said to be delocalsised.

Ionization energy.

Ionization energy is the minimum energy required to remove an electron from the outermost energy level of an atom in its gaseous state.

– The number of ionization energies an element may have is equivalent to the number of valence electrons.

– The first ionization energy is the minimum energy required to remove the first electron from the outermost energy level of an atom in its gaseous state.

– The first ionization among alkali metals decreases down the group.

Reason:

– The effective force of attraction of on the outermost electron by the positive nucleus decreases with increasing atomic size and distance from the nucleus.

– Note that the atomic radius increases down the group due to increase in the number of energy levels.

Summary on physical properties of Alkali metals.

Element	Appearance	Ease of cutting	Melting point (^oC)	Boiling point (^oC)	Electrical conductivity	Atomic radius (nm)	Ionic radius (nm)
Lithium	Silver white	Slightly hard	180	1330	Good	0.133	520
Sodium	Shiny white	Easy	98	890	Good	0.157	496
Potassium	Shiny grey	Easy	64	774	Good	0.203	419

Chemical properties of alkali metals.

Reaction with air.

– When exposed to air alkali metals react with atmospheric moisture to form the corresponding metal hydroxide and hydrogen gas.

General equation:

Metal + Water → Metal hydroxide + hydrogen gas.

– The metal hydroxide further reacts with atmospheric carbon (IV) oxide to form hydrated metal carbonate.

General equation:

Metal hydroxide + carbon (IV) oxide → Hydrated metal carbonate.

Examples:

Lithium

With moisture: 2Li_(s) + 2H₂O_(l) → 2LiOH_(aq) + H_2(g);

Then with carbon (IV) oxide:2LiOH_(aq) + CO_2(g) → Li₂CO₃.H₂O_(s);

Sodium

With moisture: 2Na_(s) + 2H₂O_(l) → 2NaOH_(aq) + H_2(g);

Then with carbon (IV) oxide: 2NaOH_(aq) + CO_2(g) → Na₂CO₃.H₂O_(s);

Potassium

With moisture: 2K_(s) + 2H₂O_(l) → 2KOH_(aq) + H_2(g);

Then with carbon (IV) oxide: 2KOH_(aq) + CO_2(g) → K₂CO₃.H₂O_(s);

Burning in air.

– Alkali metals burn in air with characteristic flame colours to form corresponding metal oxides.

Examples:

– Burns in air to form lithium oxide as the only product.

Equation:

4Li_(s) + O_2(g) → 2Li₂O_(s)

– Burns in air to with a yellow flame to form sodium oxide as the only product.

Equation:

4Na_(s) + O_2(g) → 2Na₂O_(s)

Note:

– When burned in air enriched with oxygen or pure oxygen sodium burns with a yellow flame to fom sodium peroxide (instead of sodium oxide).

Equation: 2Na_(s) + O_2(g) → Na₂O_2(s)

– Burns in air with a lilac flame to form potassium oxide as the only product.

Equation:

4K_(s) + O_2(g) → 2K₂O_(s)

Reaction with water.

– Alkali metals react with water to form the corresponding hydroxides and hydrogen gas.

Examples:

Procedure:
– A small piece of potassium metal is cut and dropped into a trough containing water;

– The resultant solution is tested with litmus paper;

Diagram of apparatus:

Observations and explanations:

– The metal floats on the water surface; because it is less dense than water;

– A hissing sound is produced; due to production of hydrogen gas;

– It explosively melts into a silvery ball then disappears because reaction between water and potassium is exothermic (produces heat). The resultant heat melts the potassium due to its low melting point.

– It darts on the surface; due to propulsion by hydrogen;

– The metal bursts into a lilac flame; because hydrogen explodes into a flame which then burns the small quantities potassium vapour produced during the reaction;

– The resultant solution turns blue; because potassium hydroxide solution formed is a strong base;

(b). Reaction equations.

Equation I

2K_(s) + 2H₂O_(l) → 2KOH_(aq) + H_2(g);

Equation II

4K_(s) + O_{2 (g)} → 2K₂O_(s);

Equation III:

K₂O_(s) + H₂O_(l) → 2KOH_(aq)

Effect of resultant solution on litmus paper;

– Litmus paper turns blue; sodium hydroxide formed is highly soluble in water; releasing a large number of hydroxyl ions which result into alkaline conditions // high pH;

Procedure:
– A small piece of sodium metal is cut and dropped into a trough containing water;

– The resultant solution is tested with litmus paper;

Diagram of apparatus:

Observations and explanations:

– The metal floats on the water surface; because it is less dense than water;

– A hissing sound is produced; due to production of hydrogen gas;

– It vigorously melts into a silvery ball then disappears because reaction between water and sodium is exothermic (produces heat). The resultant heat melts the sodium due to its low melting point.

– It darts on the surface; due to propulsion by hydrogen;

– The metal may burst into a golden yellow flame; because hydrogen may explode into a flame which then burns the sodium;

– The resultant solution turns blue; because sodium hydroxide solution formed is a strong base;

(b). Reaction equations.

Equation I

2Na_(s) + 2H₂O_(l) → 2NaOH_(aq) + H_2(g);

Equation II

4Na_(s) + O_{2 (g)} → 2Na₂O_(s);

Equation III:

Na₂O_(s) + H₂O_(l) → 2NaOH_(aq)

Effect of resultant solution on litmus paper;

– Litmus paper turns blue; sodium hydroxide formed is highly soluble in water; releasing a large number of hydroxyl ions which result into alkaline conditions // high pH;

Procedure:
– A small piece of lithium metal is cut and dropped into a trough containing water;

– The resultant solution is tested with litmus paper;

Diagram of apparatus:

Observations and explanations:

– The metal floats on the water surface; because it is less dense than water;

– A hissing sound is produced; due to production of hydrogen gas;

– It reacts less vigorously than sodium and des not melt since the melting point of lithium is relatively higher.

– It darts on the surface; due to propulsion by hydrogen;

– The gas does not ignite spontaneously;

– The resultant solution turns blue; because lithium hydroxide solution formed is a strong base;

(b). Reaction equations.

Equation I

2Li_(s) + 2H₂O_(l) → 2LiOH_(aq) + H_2(g);

Equation II

4Li_(s) + O_{2 (g)} → 2Li₂O_(s);

Equation III:

Li₂O_(s) + H₂O_(l) → 2LiOH_(aq)

Effect of resultant solution on litmus paper;

– Litmus paper turns blue; sodium hydroxide formed is highly soluble in water; releasing a large number of hydroxyl ions which result into alkaline conditions // high pH;

Summary on reaction rate.

Metal	Reaction rate
Lithium	Vigorous
Sodium	More vigorous
Potassium	Explosive

Reaction with chlorine

– All alkali metals react with chlorine to form corresponding metal chlorides.

General procedure:

– A small piece of the alkali metal is cut and placed in a deflagrating spoon;

– It is then warmed and quickly lowered into a gas jar containing chlorine.

– The experiment should be done in a working fume chamber or in the open

Reason: Chlorine gas is poisonous.

– This experiment should not be attempted in the laboratory with potassium

Reason: the reaction is too explosive and very dangerous.

Observations:

Sodium:

– The metal bursts into a yellow flame.

– White fumes of sodium chloride are formed.

Equation: 2Na_(s) + Cl_2(g) → 2NaCl_(s)

Lithium:

– The metal reacts less vigorously than sodium without bursting into a flame.

– White fumes of lithium chloride are formed.

Equation: 2Li_(s) + Cl_2(g) → 2LiCl_(s)

Potassium:

– The metal bursts into a lilac flame.

– White fumes of potassium chloride are formed.

Equation: 2K_(s) + Cl_2(g) → 2KCl_(s)

The general trend in reactivity of alkali metals.

– The reactivity of alkali metals increase down the group.

Reason:

– Alkali metals react by losing the single valence electron. The ease of loss of the valence electron increases down the group due to decrease in attraction of the valence electron towards the positive nucleus. This in turn is due to the increase in atomic radius down the group as a result of increase in number of energy levels.

Similarity of ions and formulae of some compounds of Alkali metals.

– Alkali metals have a similar charge on their ions since they all have a single valence electron.

– Thus they have the same general formula in their ions and compounds.

Examples.

Alkali metal ion	Hydroxide (OH^–)	Oxides (O^2-)	Chloride (Cl^–)	Sulphates (SO₄^2-)
Li⁺
Na⁺
K⁺

Note:

– Alkali metals are usually not found as free elements; but rather in their combined states in the earths crust.

Reason:

– They have high reactivities.

Uses of Alkali metals and their compounds.

Sodium is used in the manufacture of sodium cyanide for gold extraction.
Lithium is used in the manufacture of special high strength glasses ad ceramics.
Lithium compounds are used in the manufacture of dry cells for use in mobile phones, laptops, stopwatches, and zero emission electric vehicles.
A molten mixture of sodium and potassium is used as a coolant in nuclear reactors.
Sodium vapour is used to produce the yellow glow in street ad advertisement lights.
Molten sodium is used as a reducing agent in the extraction of titanium.

Equation: Na_(l) + TiCl_4(g) → Ti_(s)+ 4NaCl_(l)

Sodium chloride is used as a food additive.
A mixture of sodium hydroxide (caustic soda) and carbon disulphide is used in the manufacture of artificial silk called rayon.

The Alkaline Earth metals.

– Are the elements with two valence electrons and hence in group II of the periodic table.

– All are metallic in nature.

– The members of the family in order down the group is as follows:

Beryllium
Magnesium
Calcium
Strontium
Barium
Radium

Electron arrangement of the first three alkaline earth metals.

Elements	Electron arrangement
Beryllium	2.2
Magnesium	2.8.2
Calcium	2.8.8.2

Diagram: Part of periodic table showing the alkaline earth metals

Dot and cross diagrams for the first three Alkaline Earth metals

Gradation in properties of alkali metals.

Atomic and ionic radius.

Atomic radius:

– Refers to the distance between the centre of the nucleus of an atom and the outermost energy level occupied by electron(s)

Ionic radius:

– Refers to the distance between the centre of the nucleus of an ion and the outermost energy level occupied by electron(s)

Trend:

– The ionic radius and the atomic radius of alkaline earth metals increase down the group.

Reason:

– There is an increase in the number of energy levels down the group from Beryllium to radium.

Illustration:

– Beryllium (2.2) has only 2 energy levels; Magnesium (2.8.2) has 3 energy levels while calcium (2.8.8.2) has 4 energy levels.

– Thus the outermost electron in calcium is further from the nucleus than the outermost electron in magnesium and beryllium.

Atomic and ionic radius of the same element.

– For the same alkaline earth metal the atomic radius is larger than the ionic radius.

Reason:

– Alkaline earth metals form ions by losing the valence electrons, leading to the loss of an entire outermost energy level. Thus the atoms have more energy levels than the corresponding ion hence a larger radius in the atom than I the ion.

Illustration

– Calcium atom has electron arrangement of 2.8.8.2 hence 4 energy levels.

– During ion formation it reacts by losing the 2 valence electrons to acquire a new electron arrangement of 2.8.8 hence 3 energy levels.

– Thus the ion has a smaller radius than the atom.

Diagrammatically: calcium atom and calcium ion

Summary: changes in atomic and ionic radius among alkali metals.

Element	Symbol	Atomic number	Electron arrangement	Atomic radius (nm)	Ionic radius (nm)
Beryllium	Be	4	2.2	0.089	0.031
Magnesium	Mg	12	2.8.2	0.136	0.065
Calcium	Ca	20	2.8.8.2	0.174	0.099

Physical properties of alkali metals.

Appearance.

– Alkaline earth metals acquire a metallic luster when polished. This refers to a shiny appearance on the cut surface.

– They however lose this metallic luster when exposed to air due to oxidation.

Note:

– The purpose of polishing alkaline earth metals before using them in experiment is to remove the oxide coating that slows down and prevents them from reacting.

Ease of cutting.

– Magnesium is hard to cut with a knife; but is however malleable and ductile.

– Calcium cannot also be cut with a knife because it is brittle.

Note:

Malleability:
– Refers to the ability of a material to be hammered into sheets.

Example: Iron sheets are possible to be made because iron metal is malleable.

Ductility:

– The ability of a material to be rolled into wires.

Example: electric cables are made of aluminium because aluminium metal is ductile.

Brittle:

– Refers to a substance which is hard and likely to break.

They have relatively high melting and boiling points in comparison to alkali metals.

Reason: they have relatively stronger metallic bonds (than alkali metals).

– The melting and boiling points decrease down the group.

Reason: –

– The size of the atoms increase down the group due to increasing number of energy levels. As the atomic radius increase the force of attraction between the positive nucleus and the delocalized electrons decrease. This leads to a decrease in the strength of the metallic bonds (down the group).

Electrical conductivity.

– Alkaline earth metals are good conductors of heat and electricity.

Reason: they have delocalized electrons in the outermost energy level.

– The electrical conductivity is similar for all the alkaline earth metals.

Reason: all alkaline earth metals have the same number of delocalized electron (two valence electrons) in the outermost energy level.

Ionization energy.

Ionization energy is the minimum energy required to remove an electron from the outermost energy level of an atom in its gaseous state.

– The number of ionization energies an element may have is equivalent to the number of valence electrons.

– Thus alkaline earth metals have two ionization energies.

– The first ionization energy is the minimum energy required to remove the first electron from the outermost energy level of an atom in its gaseous state.

– The second ionization energy is the minimum energy required to remove the second electron from the outermost energy level of an atom in its gaseous state.

– The first and second ionization energies among alkaline earth metals decreases down the group.

Reason:

– The effective force of attraction of on the outermost electron by the positive nucleus decreases with increasing atomic size and distance from the nucleus.

– Note that the atomic radius increases down the group due to increase in the number of energy levels.

Variation between 1^st and 2^nd Ionization energies.

– The first ionization energy is always lower than the second ionization energy for the same element.

Reason:

– After losing the first electron from an atom, the overall positive charge holds the remaining electrons more firmly. Thus removing a second electron from the ion requires more energy than the first electron

Note:

– The third ionization energy will also be higher tha the second ionization energy for the same reason.

Example:
Magnesium.

First ionization energy: Mg_(g)→ Mg⁺_(g) + e- (1^stE = 736 kJ per mole)
Second ionization energy: Mg⁺_(g)→ Mg²⁺_(g) + e- (2^ndE = 1450 kJ per mole)

Summary on physical properties of alkaline earth metals.

Element	Atomic number	Melting point (^oC)	Boiling point (^oC)	Atomic radius (nm)	1^st I.E (kJmol-1)	2^nd I.E (kJmol-1)
Beryllium	4	1280	2450	0.089	900	1800
Magnesium	12	650	1110	0.136	736	1450
Calcium	20	850	1140	0.174	590	1150
Strontium	38	789	1330	0.210	550	1060
Barium	56	725	1140	0.220	503	970

Chemical properties of Alkaline earth metals.

Burning alkaline earth metals in air.

– Alkaline earth metals react burn in air to form corresponding oxides.

– More reactive alkaline earth metals may also react with atmospheric nitrogen to form corresponding nitrides.

Examples:

– Burns in air with a blinding brilliant flame forming a white solid.

– The white solid is a mixture of magnesium oxide ad magnesium nitride.

Equations:

Reacting with oxygen: 2Mg_(s) + O_2(g) → 2MgO_(s)

Reacting with nitrogen: 3Mg_(s) + N_2(g) → Mg₃N_2(s)

– Burns in air with a faint orange flame forming a white solid.

– The white solid is a mixture of calcium oxide ad calcium nitride.

Equations:

Reacting with oxygen: 2Ca_(s) + O_2(g) → 2CaO_(s)

Reacting with nitrogen: 3Ca_(s) + N_2(g) → Ca₃N_2(s)

Note:

– The trend in the reactivity of alkaline earth metals when burning in air is not clear; due to the oxide coating on the calcium that tends to slow down the reaction of calcium in air.

– For this reason it is important to polish the surfaces of alkaline earth metals before using them in experiments.

Reaction of alkaline earth metals with cold water.

– Alkaline earth metals react slowly with cold water to form corresponding hydroxides and hydrogen gas.

Examples

Magnesium:

– Reacts slowly with water to form magnesium hydroxide and hydrogen gas.

– The reaction is very slow and the amount f hydrogen gas evolved is very low hence the hydrogen gas bubbles stick on the surface of the metal.

– The magnesium hydroxide formed dissolves slightly in water to form magnesium hydroxide.

– Thus the resultant solution is slightly alkaline.

Equation:

Mg_(s) + 2H₂O_(l) → Mg(OH)_2(aq) + H_2(g).

Calcium:

Diagram of apparatus:

Observations and explanations:

– Calcium sinks to the bottom of the beaker; because it is denser than water;

– Slow effervescence of a colourless gas; due to slow evolution of hydrogen gas;

– Soapy solution formed; due to formation of alkaline calcium hydroxide;

– A white suspension is formed; because calcium hydroxide is slightly soluble in water;

Reaction equation:

Ca_(s) + H₂O_(l) → Ca (OH)_2(aq) + H_{2 (g)};

Effect of resultant solution on litmus paper;

– Litmus paper slowly turns blue; calcium hydroxide formed is slightly soluble in water; releasing a small number of hydroxyl ions which result into alkaline conditions // high pH;

Reaction with steam.

– Alkaline earth metals react with steam to produce corresponding metal oxide and hydrogen gas.

– The reactivity with stem is faster ad more vigorous for each alkaline earth metals as compared to reaction with cold water.

Examples:

Magnesium

Procedure:
– A small amount of wet sand is put at the bottom of a boiling tube;

– A small piece of magnesium ribbon is cleaned and put in the middle of the combustion tube;

– The wet sand is heated gently first then the magnesium ribbon is heated strongly until it glows.

Reason:

– To generate steam that drives out the air that would otherwise react with the magnesium (preventing reaction with steam)

– The delivery tube is removed from the water before heating stops.

Reason:

To prevent sucking back (of the gas) as the apparatus cools

– The gas produced is tested using a burning splint.

Diagram of apparatus:

Observations and explanations.

– Magnesium burns with a white blinding flame;

– Grey solid (magnesium) forms a white solid; due to formation of magnesium oxide;

– Evolution of a colourless gas that produces a pop sound when exposed to a burning splint; confirming it is hydrogen;

Reaction equation.

Magnesium + Steam → Magnesium oxide + Hydrogen gas;

Mg_(s) + H₂O_(g) → MgO_(s) + H_2(g);

Calcium

– Reaction between calcium and steam would produce calcium oxide and hydrogen gas.

– However the reaction is too explosive to be done under laboratory conditions.

Reaction with chlorine.

– Alkaline earth metals react with chlorine to form corresponding chlorides as the only products.

Condition: presence of heat, hence the metal must be heated first.

Precaution: reaction should be done in a working fume chamber because chlorine gas is poisonous).

Examples:

Procedure:
– A piece of burning magnesium is lowered into a gas jar containing chlorine.

Observations:
– The metal continues to burn with a brilliant white flame.

– They grey solid forms a white powder.

Explanation.

– Reaction between magnesium and chlorine is exothermic.

– The heat produced keeps the metal burning; thus facilitates the reaction between magnesium and chlorine to form magnesium chloride, which is the white powder.

Equation:
Mg_(s) + Cl_2(g) → MgCl_2(s)

Procedure:
– A piece of burning calcium is lowered into a gas jar containing chlorine.

Observations and explanations.

– The metal burns shortly with an orange flame but soon smolders off.

– There is no steady reaction between calcium and chlorine.

Reason:

– When calcium is heated a coating of the metal oxide is formed first which prevents furtheer reaction between the metal and chlorine.

– However under suitable conditions calcium reacts with chlorine to form a white powder of calcium chloride.

Equation:
Ca_(s) + Cl_2(g) → CaCl_2(s)

Reaction of alkaline earth metals and dilute acids

– Generally alkaline earth metals react with dilute acids to form salts and hydrogen gas.

Examples:

With hydrochloric acid.
- Beryllium:

– When a piece of beryllium is dropped into a beaker containing hydrochloric acid, there is effervescence of a colourless gas.