Chapter 02: Chemical Bonding

SYLLABUS OVERVIEW

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, $MgCl_2$, CaO.
Characteristic properties of electrovalent compounds - state of existence, melting and boiling points, conductivity, ionisation in solution, dissociation in solution and in molten state.
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. Comparison of Electrovalent and Covalent compounds.
Definition of Coordinate Bond: The lone pair effect of the oxygen atom in water and the nitrogen atom in ammonia to explain the formation of $H_3O^+$ and $OH^-$ ions in water and $NH_4^+$ ion. The meaning of lone pair.

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

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

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:

The transfer of one or more electrons from one atom to the other to form an electrovalent (or an ionic) bond.
Sharing of one, two or three pairs of electrons between two atoms to form a covalent (or a molecular) bond.
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: $Cl + e^- \rightarrow Cl^-$)

Electrovalent (or Ionic) Bond: The chemical bond formed due to the electrostatic force of attraction between a cation and an anion is called an electrovalent bond. The number of electrons that an atom of an element loses or gains is called its electrovalency.

Note: Since the electrostatic force of attraction between opposite charges is much higher, it makes the ionic compounds stable.

2.2.1 Conditions for the formation of an electrovalent bond

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.
High electron affinity: If the electron affinity value is high, anion will be formed easily.
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.

[Illustration: Orbit structure of NaCl]

AI Prompt: A clear scientific diagram showing the orbit structure of electrovalent bonding in sodium chloride. Left side: A Sodium atom (2,8,1) with an arrow pointing from its outermost electron towards a Chlorine atom (2,8,7). Right side (below): A Sodium ion (Na+) with 2,8 shells, next to a Chloride ion (Cl-) with 2,8,8 shells. Label "Coulomb forces" between the ions. Use distinct symbols like crosses and dots for the electrons of Na and Cl respectively.

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:]$^{2-}$

REDOX PROCESS IN IONIC BONDING:

In the formation of an electrovalent bond, the electropositive atom undergoes oxidation (loss of electrons), while the electronegative atom undergoes reduction (gain of electrons). Oxidation and reduction always occur simultaneously.

Example: $2Na + Cl_2 \rightarrow 2Na^+ + 2Cl^-$

Oxidation half: $2Na \rightarrow 2Na^+ + 2e^-$
Reduction half: $Cl_2 + 2e^- \rightarrow 2Cl^-$

An oxidising agent is an acceptor of electron(s) and a reducing agent is a donor of electron(s).

Intext Questions - Electrovalent Bonding

Q3 What are the conditions for the formation of an electrovalent bond?

Answer: 1. Low ionisation potential of the metallic atom. 2. High electron affinity of the non-metallic atom. 3. Large electronegativity difference between the combining atoms.

Q7 In the formation of the compound $XY_2$ an atom X gives one electron to each Y atom. What is the nature of bond in $XY_2$? Draw the electron dot structure.

Answer: The bond is electrovalent (ionic) because electrons are completely transferred from atom X to atom Y.
Electron dot structure: X: + 2 Y: $\rightarrow X^{2+} + 2[:\ddot{Y}:]^-$

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 $C_2H_2$).

Covalency: The number of electrons of an atom taking part in the formation of shared pairs. Thus, the covalency of hydrogen is 1, oxygen 2, nitrogen 3 and carbon 4.

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 $\delta+$.)

Note: A molecule that has both slight positive and slight negative charge is called a Dipole molecule. The more the electronegativity difference, the more polar the nature.
- Same EN = Non-polar
- Slightly different EN = Polar covalent
- Large EN difference = Ionic

2.3.1 Conditions for the formation of a covalent bond

Both atoms should have four or more electrons in their outermost shells (exceptions: H, Be, B, Al).
Both atoms should have high electronegativity.
Both atoms should have high electron affinity.
Both atoms should have high ionization energy.
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.

[Illustration: Structures of Water & Ammonia]

AI Prompt: Side-by-side electron dot and orbit structures for Water (H2O) and Ammonia (NH3). For Water: Central Oxygen atom overlapping with two Hydrogen atoms at an angle. Show 2 shared pairs and 2 lone pairs on Oxygen. For Ammonia: Central Nitrogen atom overlapping with three Hydrogen atoms. Show 3 shared pairs and 1 lone pair at the top of Nitrogen. Use dots and crosses for electrons.

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.

Intext Questions - Covalent Bonding

Q4 What is the difference between: (a) ionic compounds and polar covalent compounds (c) a polar covalent compound and a non-polar covalent compound?

Answer:
(a) Ionic compounds have complete transfer of electrons forming true distinct ions ($Na^+, Cl^-$). Polar covalent compounds share electrons unequally, creating only partial fractional charges ($\delta+, \delta-$).
(c) Polar covalent molecules have unequal sharing of electrons due to differing electronegativities, causing a dipole moment. Non-polar covalent molecules have equal sharing of electrons resulting in zero dipole moment.

Q13 Potassium chloride is an electrovalent compound, while hydrogen chloride is a covalent compound. But, both conduct electricity in their aqueous solutions. Explain.

Answer: KCl is an ionic compound that dissociates into pre-existing mobile ions ($K^+$ and $Cl^-$) in water. HCl is a polar covalent compound, but it chemically reacts and ionises in water ($HCl + H_2O \rightarrow H_3O^+ + Cl^-$) to also produce mobile ions. Both provide mobile ions for conduction.

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 ($H_3O^+$).

Lone Pair: A pair of electrons which is not shared with any other atom. It is provided to the other atom for the formation of a coordinate bond.

The atom which provides the electron pair is the DONOR. The atom/ion sharing it is the ACCEPTOR.

A coordinate bond has properties of both covalent and ionic bonds. It is also called a co-ionic bond.

2.6.1 Conditions for Coordinate Bond

One of the two atoms must have at least one lone pair of electrons (e.g., $NH_3, H_2O$).
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.

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

Special Case: Ammonium Chloride ($NH_4Cl$)

When $NH_4Cl$ is formed, the cation $NH_4^+$ (having 3 covalent and 1 coordinate bond) and anion $Cl^-$ are attracted towards each other by ionic bonds. Thus, ammonium chloride is a prime example of a compound having ALL THREE types of bonds: covalent, coordinate, and ionic.

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:

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

CLASS REVIEW EXERCISE

Q1 State the type of bonding in the following molecules: (a) water, (b) calcium oxide, (c) hydroxyl ion, (d) methane, (e) ammonium ion, (f) ammonium chloride.

Answer: (a) Polar Covalent (b) Electrovalent (c) Covalent (d) Non-polar Covalent (e) Covalent and Coordinate (f) Electrovalent, Covalent, and Coordinate.

Q2 Draw an electron dot diagram to show the structure of: (i) Hydronium ion, (ii) Ammonium ion, (iii) Hydroxyl ion.

Answer: (i) Hydronium ($H_3O^+$): Central O sharing 2 pairs with 2 H atoms, and 1 lone pair shared with an $H^+$. (ii) Ammonium ($NH_4^+$): Central N sharing 3 pairs with 3 H atoms, and 1 lone pair shared with an $H^+$. (iii) Hydroxyl ($OH^-$): O sharing 1 pair with H, retaining 3 lone pairs, gaining an electron.