CHEMICAL BONDING

IMPORTANT DEFINITIONS:

1. A charged particle that form from an atom (or a group of atoms) by the loss or

Gain of electrons is called an Ion.

2. A positively charged ion formed when an atom loses electron(s) is called Cation.

3. A negatively charged ion formed when an atom gains electron(s) is called Anion.

4. A simple ion is an ion formed from single atom, e.g., Na+. A polyatomic ion is an ion

Containing more than one kind of atoms, e.g. NH4 + or SO42-ion.

5. The strong electrostatic force of attraction between positive and negative ions in an

ionic compound is called Ionic Bond. e.g. NaCl,MgCl2 etc.

6. The bond formed by the sharing of electrons between two non-metal atoms is called

Covalent Bond. e.g.- CH4 , CO2 etc.

7. A molecule is a small group of atoms held together by covalent bonds.

8. The force of attraction between positive metal ions and the ‘sea of delocalized

electrons’ is called Metallic Bond.

Test Yourself:

Describe the formation of ions by electron loss/gain in order to obtain the

electronic structure of a noble gas

1.1 The Stable Noble Gas Structure:

1. The elements, helium (He), Neon (Ne), argon(Ar), Krypton (Kr), xenon (Xe) and radon (Rn)

are known as noble gases. The electronic structures (Electronic configuration) of helium,

neon and argon are shown below.

2. Helium has two valence electrons (a duplet electronic structure). All other noble gases

have eight valence electrons (an octet electronic structure).

3. An atom is stable if it has a full outer shell, that is, a duplet or octet electronic structure.

Only the noble gas atoms have this stable electronic structure.

4. The atoms of all other elements have incomplete outer shells. Hence, the atoms of these

elements will donate, accept or share electrons to achieve the stable noble gas structure

(duplet or octet configuration).

Atoms gain or lose electrons to obtain noble gases configuration.

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1.2 Formation of Positive Ions:

--------------------------------------------

1. All atoms are electrically neutral because they have equal numbers of protons and

electrons.

2. An ion is a charged particle formed when an atom loses or gains electrons.

3. Atoms of metals tend to lose valence electrons to form positive ions or cations.

4. The Positive ion contains more protons than electrons.

Example:

Formation of Sodium Ion, Na+

(a) Only the valence electron is removed from the outermost shell.

(b) The Na+ ion contains eight valence electrons (noble gas Structure).

(c) The Na+ ion carries one positive charge because it has 11 protons and only 10

Electrons.

1.3 Formation of Negative Ions:

1. Atoms of non-metals tend to gain electrons to form negative ions or anions.

2. A negative ion contains more electrons than protons.

Example:

Formation of Chloride ion, Cl-

(a) An electron is added to the outermost shell.

(b) The Cl- ion has the same electronic structure as argon.

(c) The Cl- ion carries one negative charge because it has 18 electrons and only 17

protons.

Tip for Students:

When drawing the electronic structure of an ion,the charge must be put at the top right

hand corner. You will not be given any marks if you put the charge in the nucleus.

Example 1.1:

Write down the formulae for the ions of the following elements.

Element Li Mg O Be P 1

Number of Protons

3 12 8 4 15 9

Number of Neutrons

4 12 8 4 16 9

Number of Electrons

3 10 10 2 18 10

Solution:

Ignore the number of neutrons. Calculate the total charge in the ion.

Element Li Mg O Be P F

Number of Protons

3 12 8 4 15 9

Number of Electrons

3 10 10 2 18 10

Total Charge

0 +2 -2 +2 -3 -1

Formulae Li Mg2+ O2- Be2+ P3- F-

1.4 Ionic Bond – Electron Transfer:

1. An ionic bond is formed when a metal reacts with a non-metal to form a compound.

The compound formed is called the ionic compound.

2. Ionic bonds are formed by the transfer of electrons from the metal atoms to the

non-metal atoms.

Example:

(a) Sodium (Metal) reacts with chlorine (non-metal) by the transfer of electrons to form

sodium chloride (ionic compound).

(b) There are three parts involved in the formation of sodium chloride.

Na Na+ +e-

(2,8,1) (2,8)

Cl + e- Cl-

(2,8,7)

3. The Na+ ions and Cl- ions are attracted to each other

by strong electrostatic attraction to form sodium

chloride.

This electrostatic attraction between positive and

negative ions is called the ionic bond. An ionic bond

may also be known as an electrovalent bond.

Na+ + Cl- NaCl

The formation of an ionic bond can be represented by a ‘dot and cross’ diagram.

3. An ionic bond can also be formed between a simple ion and a polyatomic ion, or between

polyatomic ions. For example, ionic bonds can be formed in magnesium sulfate (Mg2 + SO42-)

and ammonium nitrate (NH4+ + NO3

-).

4. (a) The chemical formula of an ionic compound is constructed by balancing the charges

on the Positive ions with those on the negative ions.

(b)All the positive charges must equal all the negative charges in an ionic compound.

(c) The Chemical formula of an ionic compound formed between the cation, Ma+ and the

anion, Xb - is:

2.Each chlorine atom gains an electron to form a

negatively charged Cl-ion.

1: Each sodium atom loses its single valence electron to

form a positively charged Na ‘Ion.

Ma+ Xb - Mb Xa

Example:

Fe3+ O2- Fe2O3

5. The formulae of some ionic compounds are given in the table below.

Example 1.2:

The formula of bismuth oxide is Bi2 O3 and calcium phosphate is Ca3(PO4)2. What is the charge on

(a) bismuth ion?

(b) phosphate ion?

Solution:

(a) Let the charge on bismuth ion be �. (b) Let the charge on phosphate ion be �.

The charge on oxygen ion,O2- = -2 The charge on Ca2+ = +2

2� + �(−�) = � 3(+2) + 2 � = 0

� = +� � = −�

Cation Anion Formula of ionic Compound

Na+ Cl- NaCl

Na+ OH- NaHO

K+ O2 K2O

Na+ CO2 Na2 CO3

Cation Anion Formula of ionic Compound

Mg2+ NO-3 Mg(NO3)2

Ca2 OH- Ca(OH)2

Mg2+ SO42- MgSO4

Al3+ No3- Al(NO3)3

Analysis:

An ionic compound is a neutral substance. ∴ Total positive charge = total negative charge Thus, for an ionic compound with the formula, (An+)x (B

m-)y : x(+n) + y(-m) = 0

∴ Formula of bismuth ion : Bi3+ ∴ Formula of phosphate ion: PO43-

Analysis:

• Recognise that ionic compound contain a giant lattice in which the ions are held by

electrostatic attraction.

• Deduce the formula and names of ionic compounds from their lattice structures.

1.5 Structure of ionic Compounds:

Ionic compounds form giant ionic structures because

• the positive ions and negative ions are very strongly attracted to one another,

• the ions are arranged to form a giant ionic lattice or structure.

Example:

Sodium chloride is made up of Na+ and Cl- ions . In the solid state, these ions are very

strongly attracted to one another.

Analysis:

• Relate the physical properties of ionic compounds to their lattice structure

1.6 Properties of ionic compounds:

Some of the physical properties of simple ionic substances are shown in the table below.

Property

Solubility: Ionic compounds are usually sol;uble in water, but insoluble in organic solvents like ethanol, petrol and turpentine. Ionic compounds such as silver chloride and calcium carbonate are

Explanation

Water molecules can separate the positive ions From the negative ions, Causing them to dissolve.

insoluble in water.

Volatility: Ionic compounds have high melting and Boiling points.

• The forces of attraction between positive and negative ions are very strong. • A large amount of heart energy is needed to overcome these strong bonds during melting or boiling.

Electrical conductivity: Ionic compounds are non-conductors of electricity when molten or in aqueous solution.

• In the solid state, the ions are not free to move about. • In the molten state or in aqueous solution, the ions can move freely to conduct electricity.

Tip for Students:

it is the ions (and not electrons) that conduct electricity in molten or aqueous ionic

compounds. You will get no mark if you say sodium chloride conducts electricity when

molten because ‘the ions and the electrons can move’: The incorrect use of the word

‘electrons’ negates the mark for using the word ‘ions’.

Test yourself:

• Describe the formulation of a covalent bond by the sharing of a pair of electrons

• Deduce the arrangement of electrons in covalent molecules

1.7 Covalent Bond – Sharing of Electrons:

1. A covalent bond is the chemical bond formed by the sharing of electrons between two

non-metal atoms. After bonding, each atom attains the electrons the electronic structure

of a noble gas.

2. When atoms combine by sharing electrons, molecules are formed. A molecule is a group

of two or more Atoms held together by covalent bonds.

Compounds such as hydrogen chloride (HCl) and carbon dioxide (CO2), which contain

covalent bonds, are called Covalent compounds.

3. Covalent bonds can be formed between

(a) atoms of the same element ( e. g. hydrogen, H2, and chlorine, Cl2),

(b) atoms of different elements (e. g. ammonia, NH3, and water,l H2O).

4. Covalent bonds in molecules of elements

(a) A single covalent bond or a single bond is formed when a pair of electrons are shared

between two atoms. It is represented by a single line ‘-‘ in the structural formula.

Example:

Hydrogen molecule, H2

H• + xH H•x H (or H – H)

Two hydrogen atoms One hydrogen molecule

By sharing their valence electrons each hydrogen atom achieves the stable electronic

structure by sharing their valence electrons, each hydrogen atom achieves the stable

electronic structure.

(b) A double colavent bond or a double bond is formed

when two pairs of electrons are

between two atoms. It is represented by ‘=’ in the

structure formula.

Example:

Oxygen molecule, O2

O + O O2 (or O =O)

• An oxygen atoms has six valence electrons. It needs two more electrons to have the stable octet electronic structure. • Each oxygen molecule contains a double bond.

O O O O

5. Covalent bonds in molecules of compounds

(a) Molecules made from two or more different types of atoms linked together by covalent

bonding are called molecular compounds or covalent compounds.

(b) Water, methane and carbon dioxide are example of covalent compounds.

Example:

Water molecules, H2 O

+

O + 2H H2O ( H O H )

Methane molecule, CH4

• Water is formed by the reaction of hydrogen with oxygen such that all there atoms attain noble gas structures. • Each water molecule contains two single bonds.

• Methane contains two elements carbon and hydrogen. • In a methane molecule, the carbon atoms has an octet structure while each hydrogen atoms has a duplet structure. • Each methane molecule contains four single bonds.

• Carbon dioxide is formed when carbon reacts with oxygen. • The carbon atom must share two electrons each with two oxygen atoms in order to achieve a stable octet of a electrons. • Each oxygen atoms shares two electrons and attains an octet structure.

O H H H H

O

H H

C + 4H H C H ( or H – C – H )

H H

Carbon dioxide molecules,CO2

7. An ionic compounds can contain covalent bonds,but a covalent compound cannot contain

ionic bonds.

Example:

ammonium chloride is an ionic compound that contains covalent bonds.

Tip for Students:

Compounds of metals with non-metals are usually ionic but compounds of non-metals with

non-metals (same element or different elements) are always covalent.

• Each carbon dioxide molecules contains two double bonds.

• The bonding between NH4 + and Cl- ions is an ionic bond. • The bonding between N and H atoms in NH4

+ is a covalent bond.

• Relate the physical properties of covalent substances to their structure and bonding

1.8 Properties of Covalent Substances:

1. In general,covalent substances are insoluble in water but soluble in organic solvents.

2. There are exception to ths rule. For example, hydrogen chloride (HCl), ethanol (C2H5OH )

and ethanoic acid (CH3 COOH ) are covalent substances that are soluble in water.

3. Other properties of simple covalent substances are shown in the table below.

Test Yourself:

• Compare the structure of simple molecular substances with those of giant molecular substances

in order to deduce their properties.

• Deduce the properties of Diamond, Graphite and Silicon(IV) Oxide from their bonding and

structures.

Property Low melting and boiling points

Explanation

The covalent molecules are held together by weak intermolecular forces of attraction. Only a small amount of heat energy is needed to overcome these weak forces of attraction.

Do not conduct electricity Covalent substances do not contain positive and negative ions that can move about freely. Hence, covalent substances are non-conductors of electricity (except graphite)

× During the melting of simple covalent molecules, heat energy is absorbed to break the covalent bonds between the molecules.

During the melting of simple covalent molecules, heat energy is absorbed to overcomes the attractive forces between molecules.

× Simpe covalent molecules have low melting and boiling points because covalent bonds are weak.

The covalent bonds within the molecules are strong. Simple covalent molecules have low melting and boiling points because the attractive forces between the covalent molecules are weak.

• Compare the physical properties of ionic compounds with of covalent molecules

1.9 Types of Covalent Substances:

1. Covalent substances can be divided into two categories as shown in the table below.

Simple covalent molecules Examples: hydrogen, chlorine, oxygen, water, methane and carbon dioxide

Macromolecules Example: diamond, graphite and sand (silicon dioxide)

2. Simple molecular substances:

(a) These are made up of independent molecular units. As there are on ions formed,

the forces of attraction between molecular in solid, covalent compounds like

iodine and sulfur are much weaker. They are called vander Waals’ forces and

produce a weak molecular lattice with low melting points.

(b) In covalent liquids like water, the molecules are even further apart, so the

vander Waals’forces are weaker still. In covalent gases like methane and

ammonia ,these forces are almost non-existent.

3. Giant molecular substances:

(a) Molecules with giant structures are called macromolecules.

(b) The structure, properties and uses of diamond, graphite and silicon(IV) oxide are

shown below.

4. The table below compares the physical properties of ionic compounds and covalent substances:

Diamond Graphite Silicon(IV) oxide

Giant molecular structure

• Each carbon atom has four valence electrons. • Each carbon atoms is joined to four other carbon atoms by strong covalent bonds in a tetrahedral arrangement.

• Three valence electrons in each carbon atom are used for covalent bonding and the fourth valence electron is not involved in chemical bonding. • This gives hexagonal rings of six atoms that join together to form flat layers. • These layers of carbon atoms, lie on top of each and are held together by vander Waals’ forces.

• Each silicon atom has four valence electrons and forms four covalent bonds with four oxygen atoms in a tetrahedral structure. • each oxygen atom is bonded to two silicon atoms.

Diagrammatical representation

Properties • Hard • Very high melting and boiling points • Non-conductor of electricity

• Soft and slippery • high melting and boiling points • Good conductor of electricity

• High melting and boiling points • Non-conductor of electricity

Uses • As gemstones • As tips of grinding, cutting and polishing tools

• In pencils • As a dry lubricant • Brushes for electric motor • As inert electrodes

• Main component of concrete and glass production • As an abrasive in sandblasting and in media filters for filtering water

Test yourself:

• Describe metals as a lattice of positive ions in a ‘sea of electrons’

• Define metallic bond and relate the electrical conductivity of metals to the mobility

of the electrons in the structure

1.10 Metallic Bond:

1. In a metal, the atoms are packed tightly together in a regular pattern.

2. Metals consist of a lattice of positive ions surrounded by a ‘sea of electrons’.

3. This ‘sea of electrons’ comes from the valence electrons of each atom. They are said to be

delocalised. They move freely among the metal ions like a cloud of negative charge.

4. Therefore, metallic bond is the force of attraction between positive metal ions and the ‘sea of

delocalised electrons’.

5. The lattice structure of the metallic bond is shown in the diagram below.

6. Properties of metals:

(a) Electrical conductivity:

(i) Metals are good conductors of electricity in the solid and molten states. This is due to the

‘sea of electrons’ that can move through the crystal lattice of metals.

Property Ionic compound Example: Sodium chloride

Covalent substances

Simple molecules Giant molecule Example :Iodine and Example: Diamond and carbon dioxide silicon dioxide

Volatility Non-volatile Very volatile Non-volatile

Melting and boiling points

High Low High

Solubility in water Soluble (there are some exception)

Insoluble (there are some exception)

Insoluble

solubility in organic solvents

Insoluble Soluble Insoluble

Electrical conductivity • Non-conductor in solid state • Conductor electricity in molten state or aqueous solution

Non-conductor in solid and liquid states (except graphite)

(ii) When a metal is used in an electrical circuit, electrons entering one end of the metal

cause a similar number of electrons to be displaced from the other end.

(iii) The valence electrons move from the negative terminal of the electrical circuit and

hence the metal is able to conduct an electric current.

Tips for Students:

The electrical conductivity in metals is due to the ‘moving valence electrons’.

You should not say that it is due to the ‘movement of electrons’ as it could

Include other electrons also in the metal atoms and not just the valence electrons.

TERMINAL FLOW OF ELECTRONS TERMINAL

( - ) ( + )

FREE ELECTRONS METAL ION

ELECTRICAL CONDUCTIVITY OF METALS

(b) Malleability and Ductility:

Metals are malleable (can be bent or flattened) and ductile (can be drawn into wires)

Because the layers of metal ions can slide over each other when a force is applied.

+

+

+

+ + +

+ +

+

+ +

+

+

+

+

The effect of a force being applied on a metal lattice can be seen in the diagram

below.

METAL ION VALENCE ELECTRON

APPLIED

FORCE

SLIP

PLANE

(c) Melting and boiling points:

(i) Metallic bonds are strong bonds. Hence, metals have high melting and boiling points.

(ii) There are few exceptions. For example, mercury has a low melting point, and is a

liquid at room temperature. Similarly, sodium and potassium have low melting and

and boiling points. The melting points of sodium and potassium are 98°C and

64°C respectively.

Test yourself:

• Deduce the physical and chemical properties of substances from their structures

and bonding and vice versa.

1.11 Deducing Bonding and Structure of a Substance:

We can deduce the bonding and structure of a substance from its physical properties as shown in the

in the flow chart below.

+ + +

+ + + +

+

+

+

+

+

+

+

+ +

+

+

+

+

+

+

+

+

+

+

+ + +

+ + +

Does it conduct electricity in

in the solid state?

Yes No

Does it conduct electricity in the

Liquid state?

Yes No

Is it melting point high or low

STRUCTURAL QUESTIONS AND ANSWERS

1. (a) Draw a ‘dot and cross’ diagram of the electronic structure of carbon

dioxide. Include all its electrons.

(b)Using carbon dioxide as an example, explain what is meant by covalent

bonds.

Answer:

Substances

Metal (or

graphite)

Ionic compound or

covalent molecule

Ionic compound Covalent molecule

Simple covalent Giant covalent

molecule

X

1. a) X : Electrons of Carbon

: Electrons of Oxygen

X

X C X

X X

X

b) For carbon dioxide, electrons are shared between carbon and oxygen.

The pairs of shared electrons from covalent bonds .

2.) Element Q is found in Group VI of the Periodic Table while element R is

found in Group VII of the Periodic Table. Element Q reacts with element R to

give a covalent compound.

(a) Draw a ‘dot and cross’ diagram to show the arrangement of electrons in

one molecule of the compound formed.

(b) Write the molecular formula of the compound formed.

Answer:

2. a) X : Electrons Of Q

: Electrons Of R

X X

X X

Q

X X

R R

b) QR2

3.)Explain each of the following as fully as you can.

(a)Hydrogen sulfide has a low boiling point of −60℃.

(b)Iodine does not conduct electricity.

Answer:

3.)a) Hydrogen sulfide has a simple molecular structure in which the hydrogen

sulfide molecules are held together by weak intermolecular forces . A small

amount of energy is needed to overcome the weak intermolecular forces

(holding discrete hydrogen sulfide molecules together ) .

Hence, hydrogen sulfide has a low boiling point.

b) Iodine has a simple molecular structure in which the iodine molecules

are heldtogether by weak intermolecular forces . Iodine does not

conduct electricity because of the absence of mobile charged particles

(such as ions or electrons ) .

Answer:

4.) Hydrogen exists as three isotopes: hydrogen, deuterium and tritium.

(a) Using D to represent deuterium, construct a ‘dot and cross’ diagram to

show the bonds in the oxide of deuterium.

(b) Using T to represent tritium, construct a ‘dot and cross’ diagram to show

the arrangement of electrons in a tritium molecule.

Answer:

4.) a) X : Electrons of Oxygen

: Electrons of Deuterium

X X

X X X

X

X X

D D

b)

T T

X

5.) (a) Write the formula of a chloride ion.

(b) Draw a diagram to illustrate the formation of a chloride ion from a

chlorine atom. Show the electronic structures of both particles.

Answer:

5.) a) Cl-

b)

Gains An

Electron

Cl Cl-

6.) (a) Draw a ‘dot and cross’ diagram of the electronic structure of

magnesium oxide.Include all its electrons.

(b) Using magnesium oxide as an example, explain what is meant by

ionic bonds.

17p

18n

17p

18n

Answer:

6.) a)

2+ X X 2-

X

X

X O X

X

X

Mg2+ O2-

b) In magnesium oxide, two is the electrons from a magnesium atom are

transferred to an oxygen atom. The ionic bond is the electrostatic force of attraction

between the positively – charged magnesium ion and the negatively – charged oxide

ion .

7.) Explain each of the following as fully as you can.

(a) Magnesium oxide, which has a melting point of 2800 ℃, is used as a

refractory material.

(b) Sodium chloride conducts electricity in the molten state but does not

conduct electricity in the solid state.

(c) Calcium fluoride is a brittle material.

Answer:

7.) a) Magnesium oxide has a giant ionic structure ( lattice ) in which the

positively charged magnesium ions ( Mg2+ ) and the negatively – charged

oxide ions ( O2- ) are held together by strong electrostatic attraction

Mg

( ionic bonds ) . A lage amount of energy is needed to overcome the strong

ionic bonds holding the Mg2+and O2- ions together. Hence, magnesium

oxide has a high melting point and can be used as a refractory material .

b) In the solid state, sodium chloride has a giant ionic structure ( lattice )

in which the positively – charged sodium ions and the negatively – charged

chloride ions are held together by strong electrostatic attraction.

These ions are held in fixed positions and are not fro to move about . Thus,

sodium chloride cannot conduct electricity in the solid state. In the molten

state, electrostatic attraction is slightly weaker and the ions are free to

move about (slide around one another ) . Thus, sodium chloride in the

molten state conducts electricity because of the presence of mobile ions.

c) Calcium fluoride has a giant ionic structure ( lattice ) in which the

positively – charged calcium ions and the negatively – charged fluoride

ions are held together by strong electrostatic attraction ( ionic bonds ) .

When a force is applied, a slight displacement in the lattice structure

occurs and brings similarly charged ions together . Repulsion between the

like charges fractures the lattice of calcium fluoride easily . Hence, calcium

fluoride is a material.

8.)Explain each of the following as fully as you can.

(a) Iron has a high melting point of 1538℃.

(b) Aluminium is malleable and ductile.

(c) Copper is a good conductor of electricity.

Answer:

8.) a) Iron has a giant metallic structure in which a lattice of positive ions and

delocalized electrons are bound together by strong metallic bonds . A large

amount of energy is needed to overcome the strong metallic bonds in iron .

Hence, iron has a high melting point .

b) Aluminium has a giant metallic structure in which a lattice of positive

ions and delocalized electrons are bonded by strong melting bonds . The

layers of positive ions can slide over each other without breaking the metallic

bonds . After being beaten into different shapes or drawn into wires, the

neighbouring positive ions are still held together by the ‘ Sea of electrons ’ .

c)Copper has a giant metallic structure in which a lattice of positive ions and

delocalised electrons are bound together by strong metallic bonds.

Copper is able to conduct electricity because of the presence of the sea of

mobile delocalised electrons .