What is a Chemical Bond?

 

What is a Chemical Bond?

A chemical bond is a force that holds two or more atoms together in a molecule of compound.

Atoms bond together to achieve a stable electronic configuration, usually the octet (8 electrons in their outermost shell), like the noble gases.

Why Do Atoms Form Bonds?

Atoms form chemical bonds to achieve stability by reducing their potential energy.

Main Assumptions of Kossel-Lewis Theory

(i) Atoms tend to attain octet configuration (8 electrons in valence shell).

Noble gases (except He) have stable octet configurations, which is why they are chemically inert.

(ii) Atoms can achieve octet by:

* Losing electrons → forming positive ions (cations)

* Gaining electrons → forming negative ions (anions)

* Sharing electrons → forming covalent bonds

 

Kossel’s Contribution – Ionic Bonding

Ionic bond formation involves:

* Metal atoms (with 1–3 valence electrons) lose electrons to form cations.

* Non-metal atoms (with 5–7 valence electrons) gain electrons to form anions.

* Resulting oppositely charged ions attract each other to form ionic bonds.

Example: Formation of NaCl

Na (2,8,1) → Na⁺ + e⁻

Cl (2,8,7) + e⁻ → Cl⁻

Na⁺ and Cl⁻ combine to form NaCl via electrostatic attraction.

Lewis’s Contribution – Covalent Bonding

Covalent bond formation involves:

* Atoms sharing valence electrons to attain octet/duplet configuration.

* Shared electron pair belongs to both atoms.

* Lewis represented atoms using Lewis dot symbols:

* Valence electrons shown as dots around the element symbol.

Example: Formation of H₂ molecule

Each H atom has 1 electron (needs 2 for duplet)

They share one pair of electrons: H:H

Each H attains a duplet → stable H₂ molecule

Octet Rule

Atoms tend to gain, lose, or share electrons to attain 8 electrons in their outermost shell.

·        Helium (He) is an exception (only 2 electrons – duplet rule).

Limitations of Octet Rule:

Doesn’t apply to:

* Odd-electron molecules (e.g., NO, NO₂)

* Incomplete octet (e.g., BCl₃)

* Expanded octet (e.g., PCl₅, SF₆)

 

Modes Of Chemical Combination

·        Electrovalent bond or ionic bond

·        Covalent bond

·        Co-ordinate bond

 

Electrovalent Bond / Ionic Bond

Definition

An electrovalent bond is a chemical bond formed by the complete transfer of electrons from one atom to another. It also known as an ionic bond.

* The atom that loses electrons becomes a positively charged ion (cation).

* The atom that gains electrons becomes a negatively charged ion (anion).

* These oppositely charged ions are held together by strong electrostatic force of attraction.

Example: Formation of NaCl (Sodium Chloride)

Step 1: Electron transfer

Na (2,8,1) → Na⁺ + e⁻

Cl (2,8,7) + e⁻ → Cl⁻

Step 2: Formation of ions

Na⁺ (cation), Cl⁻ (anion)

Step 3: Attraction

Na⁺ and Cl⁻ attract each other and form NaCl

Na  + Cl  → Na⁺ + Cl⁻    → NaCl

 

Energy Changes During Ionic Bond Formation

Forming an ionic bond involves multiple energy steps:

(i) Ionization Energy (IE)

Energy required to remove electron(s) from a metal atom.

(ii) Electron Affinity (EA)

Energy released when a non-metal gains electron(s).

(iii) Lattice Energy

Energy released when oppositely charged ions come together to form a solid ionic lattice.

Higher lattice energy = more stable ionic compound

Example: MgO has high lattice energy → very stable compound.

 

Conditions for Formation of Ionic Bond

To form an ionic bond, the following conditions are generally required:

(i) Low Ionization Energy of the metal

So that it can easily lose electrons and form a cation.(e.g., Na, Mg, Ca)

(ii) High Electron Affinity of the non-metal

So that it can easily gain electrons and form an anion. (e.g., Cl, O, F)

(iii) Large Electronegativity Difference

Difference in electronegativity between metal and non-metal should be greater than 1.7.

Factors Affecting Strength of Ionic Bond

(i) Charge on Ions

Greater charge = stronger bond

e.g., MgO (Mg²⁺ and O²⁻) has stronger bonding than NaCl (Na⁺ and Cl⁻)

(ii) Size of Ions

Smaller ions = stronger attraction

e.g., LiF has stronger bonding than CsI

Characteristics of Ionic Compounds

(i) Physical State

Generally exist as crystalline solids.

(ii) Melting and Boiling Points

High due to strong electrostatic forces.

(iii) Solubility

* Soluble in polar solvents like water.

* Insoluble in non-polar solvents like benzene.

(iv) Electrical Conductivity

* Good conductors in molten or aqueous state (ions are free to move).

* Poor conductors in solid state (ions are fixed).

(v) Bond Strength

Strong bonds due to strong attraction between oppositely charged ions.

 Advantages of Ionic Bonding

* Explains the formation and properties of salts.

* Helps understand reactivity of metals and non-metals.

Limitations of Ionic Bonding

* Ionic bonding does not explain the properties of compounds with covalent character.

* Cannot be formed if electronegativity difference is too small.

* Compounds may show partial covalent character (Fajans’ Rule).

 

Fajans’ Rule (Covalent Character in Ionic Compounds)

* It explains why some ionic compounds show covalent character.

e.g  ,AlCl₃ shows covalent nature despite being made of metal and non-metal.

* It is based on the concept of polarization.Example:

Polarization

Polarization is the distortion of the electron cloud of an anion by a cation.

* The greater the polarization, the greater the covalent character in the compound.

Factors affecting Covalent Character (Fajans’ Rule)

(i) Size of the Cation

Smaller cation → greater polarizing power → more covalent character

Example: Al³⁺ (small) > Na⁺ → AlCl₃ more covalent than NaCl

(ii) Charge on the Cation

Higher charge → stronger attraction to electrons → more distortion

Example: Fe³⁺ > Fe²⁺ → FeCl₃ more covalent than FeCl₂

(iii) Size of the Anion

Larger anion → loosely held electrons → easily distorted → more covalent

Example: I⁻ > Br⁻ > Cl⁻ > F⁻

So, NaI > NaBr > NaCl > NaF (in covalent character)

(iv) Charge on the Anion

Higher negative charge → more electrons to polarize → more covalent character

Example: O²⁻ > F⁻

→ MgO is more covalent than MgF₂

 

Factors Affecting Covalent Character

Factor                     Trend for More Covalent Character

Cation size             Smaller

Cation charge        Higher

Anion size               Larger

Anion charge           Higher

 

Examples Based on Fajans’ Rule

Compound   Why It Is More Covalent?          Nature

AlCl₃   Al³⁺ is small and highly charged → strong polarization       More covalent

NaCl   Na⁺ is large, Cl⁻ is small → less polarization    Mostly ionic

BeCl₂  Be²⁺ is very small → high polarizing power    Covalent

MgI₂   I⁻ is large → easily polarized       Partially covalent

FeCl₃   Fe³⁺ more polarizing than Fe²⁺   More covalent

 

Application of Fajans’ Rule

1.Predicts the degree of covalent character in ionic compounds.

2. Explains why:

* Some ionic compounds are insoluble in water.

* Some ionic compounds have low melting points.

* Transition metal halides often show covalent nature.

* AlCl₃ is covalent and exists as a dimer in vapor phase.

Crystalline solids

Crystalline solids are a type of solid material where the atoms, ions, or molecules are arranged in a highly ordered and repeating pattern extending in all three spatial dimensions.

Characteristics of Crystalline Solids:

1. Definite Geometric Shape

Due to the regular arrangement of particles, crystalline solids often form flat surfaces and sharp edges—these are called crystal faces.

2. Long-Range Order

The internal structure is consistent and periodic across large distances.

3. Sharp Melting Point

Crystalline solids melt at a specific temperature, because all bonds of the same type break at once.

4. Cleavage Planes

They tend to break along specific planes where atomic bonding is weaker.

Examples of Crystalline Solids:

* Table salt (NaCl) – Ionic crystal with cubic symmetry

* Diamond – Covalent network crystal with tetrahedral bonding

* Quartz (SiO₂) – Covalent network with repeating tetrahedra

* Copper (Cu) – Metallic crystal with good electrical conductivity