ICSE Class 10 Chemistry Chapter 02 Chemical Bonding

Chemical Bonding

Syllabus

Electrovalent, covalent and co-ordinate bonding, structures of various compounds - orbit structure and electron dot structure.

Definition of Electrovalent Bond.

Structure of Electrovalent compounds NaCl, MgCl2, CaO:

Characteristic properties of electrovalent compounds - state of existence, melting and boiling points, conductivity (heat and electricity), ionisation in solution, dissociation in solution and in the molten state to be linked with electrolysis.

Covalent Bond - definition and examples, structure of Covalent molecules on the basis of duplet and octet of electrons (example: hydrogen, chlorine, nitrogen, water, ammonia, carbon tetrachloride, methane)

Characteristic properties of Covalent compounds - state of existence, melting and boiling points, conductivity (heat and electricity), ionisation in solution.

Comparison of Electrovalent and Covalent compounds.

Definition of Coordinate Bond: The lone pair effect of the oxygen atom of the water molecule and the nitrogen atom of the ammonia molecule to explain the formation of H3O+ and OH- ions in water and NH4+ ion. The meaning of lone pair; the formation of hydronium ion and ammonium ion must be explained with the help of electron dot diagrams.

Introduction

Everything in this world wants stability so is the case with atoms. For atoms, stability means having the electron arrangement of an inert gas, i.e., octet in its outermost shell. Helium has two electrons (DUPLET) while all other inert gases, i.e., Neon, Argon, Krypton, Xenon and Radon have eight electrons (OCTET) in their outermost shell, as given in Table 2.1 below.

Table 2.1 Electronic configurations of the inert gases

It is found that the elements with their complete outermost shell do not react or are least reactive. We, therefore, conclude that the atoms having 8 electrons (or 2 electrons, Helium configuration) in their outermost shells are very stable and unreactive. Therefore, to attain stability, atoms tend to combine chemically by redistribution of electrons in the outermost shell or valence electrons so that each is left with a stable electronic configuration (duplet or octet).

Cause of chemical combination is the tendency of elements to acquire the nearest noble gas configuration in their outermost orbit and become stable.

During redistribution of electrons, a bond (force) of attraction develops between atoms, which binds them together to form molecules. This bond is known as the chemical bond.

A chemical bond may be defined as the force of attraction between any two atoms, in a molecule, to maintain stability.

There are three methods in which atoms can achieve a stable configuration.

1. The transfer of one or more electrons from one atom to the other to form an electrovalent (or an ionic) bond.

2. Sharing of one, two or three pairs of electrons between two atoms to form a covalent (or a molecular) bond.

3. When the shared electron pairs are contributed by only one of the combining atoms, the bond formed is known as coordinate (or dative) bond.

Teacher's Note

Chemical bonds are like the social connections between people - just as people bond to feel more stable and complete, atoms bond to achieve stability and lower energy states.

Electrovalent (Or Ionic) Bond

Atoms of metallic elements that have 1, 2 or 3 valence electrons can lose electron(s) to atoms of non-metallic elements, whichs have 7, 6 or 5 electrons respectively in their outermost shell and thereby forming an electrovalent compound.

After the transfer of electron(s), both the combining atoms acquire the electronic configuration of the nearest inert gas.

A metallic atom, which loses electron(s), becomes a positively charged ion and is known as a cation and a non-metallic atom, which gains electron(s), becomes a negatively charged ion and is known as an anion.

An ion is a charged particle which is formed due to the gain or the loss of one or more electrons by an atom.

A metallic element, whose one atom readily loses electron(s) to form a positively charged ion, is an electropositive element.

Na - e- - Na+ (cation)

A non-metallic element, whose atom readily accepts electron(s) to form a negatively charged ion, is an electronegative element.

Cl + e- - Cl- (anion)

The cation and the anion being oppositely charged attract each other and form a chemical bond. Since this chemical bond formation is due to the electrostatic force of attraction between a cation and an anion, it is called an electrovalent (or an ionic) bond.

Electrovalent (or ionic) compounds: The chemical compounds formed as a result of the transfer of electrons from one atom of an element to one atom of another element are called ionic (or electrovalent) compounds.

Electrovalency: The number of electrons that an atom of an element loses or gains to form a electrovalent bond is called its electrovalency.

Conditions for the formation of an electrovalent (or ionic) bond

When an ionic compound is formed, the neutral atom is changed to a cation or an anion. The formation of cations and anions depends on the following factors:

1. Low ionisation potential: If the ionisation potential of a particular atom is low, it will lose electron(s) easily, i.e., a cation is formed easily.

2. High electron affinity: If the electron affinity value is high, anion will be formed easily, i.e., a higher electron affinity value favours ionic bonding.

3. Large electronegativity difference: If the difference in the electronegativities of two elements is higher, then the transfer of electrons will be easier. Therefore, more the difference in electronegativity, more will be the ionic nature of the resulting compound.

The metals of groups 1, 2 and 13 have a tendency to lose their valence electrons. So they combine with the non-metals of Groups 15, 16 and 17, which have a tendency to gain electron(s) and form ionic bonds.

Note: Group 1 elements are most electropositive, i.e., they are metallic in nature. Their metallic nature increases down the group.

Group 17 elements are most electronegative. Fluorine is the most electronegative element. Thus, caesium fluoride CsF is the most ionic compound.

Bonds formed between metals and non-metals are ionic or electrovalent.

Why are ionic Compounds Stable?

Ionic compounds are formed by ions but there also exists a repulsive force between ions for like charges. Since the electrostatic force of attraction between opposite charges is much higher, it makes the ionic compounds stable.

Examples of electrovalent (ionic) compounds

NaCl - (Sodium chloride)

MgCl2 - (Magnesium chloride)

CaO - (Calcium oxide)

KBr - (Potassium Bromide)

CaCl2 - (Calcium chloride)

Teacher's Note

Ionic bonding is like trading tokens with classmates to get what you need - sodium gives away electrons (like tokens) to chlorine, making both chemically happier and more stable.

Structures of some electrovalent compounds

1. Sodium chloride (NaCl)

The electronic configuration of a sodium atom is 2, 8, 1. It has one electron in excess of the stable electronic configuration of the nearest noble gas, neon, (2, 8). Therefore, an atom of sodium shows a tendency to give up the electron from its outermost shell, so as to acquire a stable electronic configuration of neon.

Na - 1e- - Na+

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

atom - cation

However, after giving up one electron, the sodium atom is no more electrically neutral. It has eleven protons in its nucleus but only ten electrons revolving around it. Therefore, it has a net positive charge of +1. This positively charged atom is called sodium ion and is written as Na+ and its electronic configuration resembles that of the noble gas neon.

The electronic configuration of chlorine is 2, 8, 7. It has an electronic configuration with one electron less than that of the nearest noble gas, argon (2, 8, 8). Therefore, the chlorine atom shows a tendency to acquire an electron to attain octet in its outermost shell.

Cl + 1e- - Cl-

(2, 8, 7) - (2, 8, 8)

atom - anion

An atom of chlorine is electrically neutral, as it contains 17 protons in its nucleus and 17 electrons revolving around it. But, after acquiring an electron from the sodium atom, the chlorine atom does not remain electrically neutral. It has one electron more than the number of protons in its nucleus and therefore has charge of -1 represented as Cl- i.e. chloride ion.

Chloride ion has an octet of electrons in its outermost shell, and its electronic configuration resembles that of the noble gas argon (Fig. 2.1).

Thus, when an atom of sodium combines with an atom of chlorine (electronic configuration 2, 8, 7), one electron is transferred from the sodium atom to the chlorine atom, resulting in the formation of a sodium chloride molecule.

Thus, the ratio of magnesium to chloride ions in magnesium chloride must be 1 : 2 ; so the molecular formula of the compound magnesium chloride is MgCl2 (Fig. 2.2).

Electron dot structure of magnesium chloride

Mg - + 2[Cl-] - Mg2+ 2[Cl-]

3. Calcium oxide (CaO)

The number of valence electrons of a calcium atom (atomic number 20) is 2, and that of an oxygen atom is 6, i.e., oxygen requires 2 electrons to attain octet. In the presence of oxygen, each calcium atom loses its 2 valence electrons to the oxygen atom (Fig. 2.3). As a result, the calcium atom forms a calcium ion with charge +2 (Ca2+), and the oxygen atom forms an oxide ion with charge -2 (O2-). Since only one oxygen atom is needed to accept the 2 valence electrons donated by a calcium atom, the formula of calcium oxide is CaO and not Ca2O2.

Ca - 2e- - Ca2+

(2, 8, 8, 2) - (2, 8, 8)

atom - cation

O + 2e- - O2-

(2, 6) - (2, 8)

atom - anion

Electron dot structure of calcium oxide

Ca + O - Ca2+ [O2-]

In the formation of an electrovalent bond, the transfer of electron(s) is involved. The electropositive atom undergoes oxidation, while the electronegative atom undergoes reduction. This is known as REDOX PROCESS.

For example:

Formation of sodium chloride: Sodium chloride is formed by the combination of sodium and chlorine.

2Na + Cl2 - 2Na+ + 2Cl- (or 2 NaCl)

The reaction can be written as two half reaction:

2Na - 2Na+ + 2e- (Oxidation)

Cl2 + 2e- - 2Cl- (Reduction)

Oxidation

2Na + Cl2 - 2Na+ + 2Cl- (Redox Reaction)

Reduction

Oxidation and reduction always occur simultaneously because the electron(s) lost by the reducing agent must be gained by the oxidising agent.

For example:

Oxidised

CuO + H2 - Cu + H2O

Oxidising agent - Reducing agent

Reduced

In this reaction, hydrogen acting as a reducing agent reduces Cu(II) oxide to copper. This is a reduction reaction.

Cu2+ + 2e- - Cu (Reduction)

At the same time, hydrogen is oxidised to water by the oxidising agent Cu(II) oxide, and this is an oxidation reaction.

2H - 2e- - 2H+ (Oxidation)

Thus, the net reaction is a redox reaction.

It can be inferred from the above example that an oxidising agent is an acceptor of electron(s) and a reducing agent is a donor of electron(s).

[Also refer chapter 6, article 6.2 (viii) & (ix)]

Teacher's Note

When you trade a toy with a friend, you give up one thing and receive another - similarly, in ionic bonding, atoms give and receive electrons to both become more stable and satisfied.

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