2.1 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.

Table 2.1: Electronic configurations of the inert gases

Inert gas	Atomic No.	K	L	M	N	O	P	No. of valence electrons
He	2	2						2
Ne	10	2	8					8
Ar	18	2	8	8				8
Kr	36	2	8	18	8			8
Xe	54	2	8	18	18	8		8
Rn	86	2	8	18	32	18	8	8

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 a coordinate (or dative) bond.

2.2 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, which have 7, 6 or 5 electrons respectively in their outermost shell and thereby forming an electrovalent compound.

• A metallic atom, which loses electron(s), becomes a positively charged ion and is known as a cation. (Electropositive element: $Na - e^- \rightarrow Na^+$)
• A non-metallic atom, which gains electron(s), becomes a negatively charged ion and is known as an anion. (Electronegative element: $Cl + e^- \rightarrow Cl^-$)

2.2.1 Conditions for the formation of an electrovalent bond

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.
3. Large electronegativity difference: More the difference in electronegativity, more will be the ionic nature of the resulting compound.

Key Note: Group 1 elements are most electropositive. Group 17 elements are most electronegative (Fluorine is the most). Thus, caesium fluoride (CsF) is the most ionic compound.

2.2.2 Structures of some electrovalent compounds

1. Sodium chloride (NaCl)

Sodium ($_{11}Na$: 2, 8, 1) loses 1 electron. Chlorine ($_{17}Cl$: 2, 8, 7) gains 1 electron.

Electron dot structure of NaCl:

Na• + :Cl: → Na$^+$ + [:Cl:]$^-$ or NaCl

Table 2.2: Comparison of Sodium atom and Sodium ion

Property	Sodium atom (Na)	Sodium cation (Na$^+$)
Colour	Silvery white	Colourless
Toxicity	Poisonous	Non-poisonous
Chemical action	Very active	Inactive
Valence shell	Incomplete outermost shell	Complete outermost shell
Electrical state	Neutral	Positively charged
Existence	Combined state	Independent existence

Table 2.3: Comparison of Chlorine atom and Chloride ion

Property	Chlorine atom (Cl)	Chloride anion (Cl$^-$)
Colour	Yellowish green (as $Cl_2$ gas)	Colourless
Toxicity	Poisonous	Non-poisonous
Odour	Suffocating	Odourless
Chemical action	Very active	Inactive
Valence shell	Incomplete outermost shell	Complete outermost shell
Electrical state	Neutral	Negatively charged
Existence	Not independent	Independent

2. Magnesium chloride ($MgCl_2$)

Magnesium ($_{12}Mg$: 2, 8, 2) loses 2 electrons. It gives one electron to each of the TWO Chlorine atoms ($_{17}Cl$: 2, 8, 7).

Mg: + 2 :Cl: → Mg$^{2+}$ + 2[:Cl:]$^-$

3. Calcium oxide (CaO)

Calcium ($_{20}Ca$: 2, 8, 8, 2) loses 2 electrons. Oxygen ($_{8}O$: 2, 6) gains 2 electrons.

Ca: + :O: → Ca$^{2+}$ + [:O:]$^{2-}$

2.3 COVALENT (MOLECULAR) BOND

The chemical bond that is formed between two combining atoms by mutual sharing of one or more pairs of electrons is called a covalent bond. The atoms of non-metals usually have 5, 6 or 7 electrons (except C with 4, H with 1) in their outermost shell. Since they do not favour the loss of electrons, they complete their octet by sharing.

• Single covalent bond (-): Sharing of one pair of electrons (e.g., $Cl-Cl$, $H-Cl$, $H_2O$, $NH_3$, $CH_4$, $CCl_4$).
• Double covalent bond (=): Sharing of two pairs of electrons (e.g., $O=O$, Ethene $C_2H_4$).
• Triple covalent bond (≡): Sharing of three pairs of electrons (e.g., $N\equiv N$, Ethyne $C_2H_2$).

Polar and Non-Polar Covalent Compounds

• Non-polar covalent compounds: Shared pair of electron(s) are equally distributed. No charge separation. Molecule is symmetrical and electrically neutral. Examples: $H_2, Cl_2, O_2, CH_4, CCl_4$. (Methane is non-polar because C(2.5) and H(2.1) have nearly equal electronegativities, making the molecule symmetrical.)
• Polar covalent compounds: Shared pair of electrons are not at equal distance. This results in fractional positive ($\delta+$) and negative ($\delta-$) charges. They ionise in water. Examples: $HCl, H_2O, NH_3$. (In HCl, Cl is highly electronegative (3.0) and pulls the shared pair towards itself, becoming $\delta-$, leaving H as $\delta+$.)

2.3.1 Conditions for the formation of a covalent bond

1. Both atoms should have four or more electrons in their outermost shells (exceptions: H, Be, B, Al).
2. Both atoms should have high electronegativity.
3. Both atoms should have high electron affinity.
4. Both atoms should have high ionization energy.
5. The electronegativity difference between the combining atoms should either be zero or negligible.

Table: Decreasing Electrovalent Character (Period 3 Trends)

Group →	Group I	Group II	Group III	Group IV	Group V	Group VI	Group VII
Chloride →	NaCl	$MgCl_2$	$AlCl_3$	$SiCl_4$	$PCl_3/PCl_5$	$S_2Cl_2$	-
Bonding →	Ionic solid	Ionic solid	Partially ionic/covalent	Covalent liquid	Covalent liq/solid	Covalent liquid	-
Oxide →	$Na_2O$	MgO	$Al_2O_3$	$SiO_2$	$P_2O_5$	$SO_2/SO_3$	$Cl_2O_7$
Bonding →	Ionic Solid	Ionic Solid	Ionic Solid	Covalent Solid	Covalent Solid	Covalent Gas/Solid	Covalent Gas

2.3.2 Structures of Covalent Molecules

1. Hydrogen Molecule ($H_2$) - Non-polar: Mutual sharing of one pair of electrons. Single covalent bond [H - H].

2. Chlorine Molecule ($Cl_2$) - Non-polar: Each Cl needs 1 electron. Mutual sharing of one pair. Single covalent bond [Cl - Cl].

3. Nitrogen Molecule ($N_2$) - Non-polar: Each N needs 3 electrons. Mutual sharing of THREE pairs of electrons. Triple covalent bond [$N \equiv N$].

4. Water Molecule ($H_2O$) - Polar: Oxygen (2,6) shares two of its electrons with two different Hydrogen atoms. Two single covalent bonds. Molecule has 2 lone pairs on Oxygen.

5. Ammonia Molecule ($NH_3$) - Polar: Nitrogen (2,5) shares three of its electrons with three different Hydrogen atoms. Three single covalent bonds. Molecule has 1 lone pair on Nitrogen.

6. Carbon Tetrachloride ($CCl_4$) - Non-polar: Carbon (2,4) shares its 4 valence electrons with 4 different Chlorine atoms. Four single covalent bonds. Although C-Cl bonds are polar, the symmetrical tetrahedral shape cancels out the dipoles, making the overall molecule non-polar.

7. Methane ($CH_4$) - Non-polar: Carbon (2,4) shares its 4 valence electrons with 4 different Hydrogen atoms. Four single covalent bonds.

2.4 PROPERTIES & COMPARISON OF ELECTROVALENT AND COVALENT COMPOUNDS

Electrovalent compounds		Covalent compounds
1. Nature Constituent particles are ions. Hard solids.	Strong electrostatic forces of attraction between ions.	1. Nature Constituent particles are molecules. Gases, liquids or soft solids.	Weak forces of attraction (Vander Waals) between molecules.
2. Boiling & Melting Point Non-volatile, High BP and MP.	Large amount of energy required to break the strong ionic bonds.	2. Boiling & Melting Point Volatile, Low BP and MP.	Less energy required to break weak intermolecular forces.
3. Conductivity Do not conduct in solid state. Good conductors in fused/aqueous state.	Electrostatic forces weaken in fused/solution state. Ions become mobile.	3. Conductivity Non-conductors in solid, molten or aqueous state.	Due to the absence of free mobile ions.
4. Ionisation in solution They are electrolytes.	Water (polar) decreases electrostatic forces, resulting in free ions ($Na^+$, $Cl^-$).	4. Ionisation in solution Non-polar do not ionise. Polar ones (like HCl) ionise and act as electrolytes.	Polar covalent molecules form ions in their solutions ($HCl + H_2O \rightarrow H_3O^+ + Cl^-$).
5. Dissociation Ions dissociate in water or molten state.	Pre-existing ions separate out.	5. Dissociation Does not take place.	Covalent compounds do not have pre-existing ions.
6. Solubility Soluble in water, insoluble in organic solvents.	Like dissolves like. Water separates the ions.	6. Solubility Insoluble in water, dissolve in organic solvents.	Non-polar molecules dissolve in non-polar organic solvents.
7. Rate of reaction Rapid speed of chemical reactions.	Free ions in solutions unite very fast.	7. Rate of reaction Slow speed.	Old bonds must be broken and new bonds formed.

2.5 EFFECT OF ELECTRICITY ON COMPOUNDS

Experiment: When electric current is passed through solutions of electrovalent compounds (NaCl, $MgCl_2$, $NaOH$, $CuSO_4$), the bulb glows. This shows they are good conductors due to free mobile ions.

When passed through solutions of covalent compounds (distilled water, sugar solution, alcohol, chloroform, benzene), the bulb does not glow. They contain only molecules and no ions.

2.6 COORDINATE BOND

The bond formed between two atoms by sharing a pair of electrons, provided entirely by one of the combining atoms but shared by both, is called a coordinate bond or dative bond. e.g., Ammonium ion ($NH_4^+$), Hydronium ion ($H_3O^+$).

2.6.1 Conditions for Coordinate Bond

1. One of the two atoms must have at least one lone pair of electrons (e.g., $NH_3, H_2O$).
2. Another atom should be short of at least a lone pair of electrons (e.g., $H^+$ ion, which has 0 electrons).

2.6.2 Formation of Ammonium ion ($NH_4^+$)

In ammonia ($NH_3$), Nitrogen shares 3 electrons with 3 Hydrogen atoms and is left with one lone pair. When $NH_3$ combines with a hydrogen ion ($H^+$), which has no electrons, the lone pair on Nitrogen is shared by the $H^+$. Once formed, all four N-H bonds become identical. The linkage forms an ammonium ion having a single positive charge.

$NH_3 + H^+ \rightarrow NH_4^+$

2.6.3 Formation of Hydronium ion ($H_3O^+$) and lone pair effect of Oxygen

Water ($H_2O$) has an Oxygen atom sharing 2 electrons with 2 Hydrogens, leaving two lone pairs. Because Oxygen is highly electronegative, water is a polar molecule ($H^{\delta+} - O^{\delta-} - H^{\delta+}$).

When an acid is added to water, the polar interaction releases an $H^+$ ion from the acid. This $H^+$ spontaneously adds onto one of the lone pairs of the oxygen atom, forming the Hydronium ion (a hydrated proton).

$H_2O + H^+ \rightarrow H_3O^+$

2.6.4 Formation of Hydroxyl ion ($OH^-$) and Self-ionisation of water

The hydroxyl ion is formed when one $H^+$ is removed from water. The shared pair of electrons remains with oxygen (since it's more electronegative), giving $OH^-$ a negative charge.

$H_2O \rightarrow H^+ + OH^-$

In water, self-ionisation occurs where an $H^+$ is transferred from one water molecule to the lone pair of another, forming both ions:

$H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$

CHAPTER EXERCISE (Exam Prep & Past Board Questions)