NIOS Final Exam Focus

Lesson 4: Chemical Bonding

This chapter is highly important for exams. Focus on bond formation, ionic and covalent compounds, VSEPR shapes, hybridization, resonance, and molecular orbital theory.

1. Valence Electrons and Bond Formation

Valence electrons are electrons in the outermost shell.
They take part in bond formation.
Chemical bond forms when energy of the system decreases.
Atoms combine to achieve stable noble gas configuration.

2. Types of Bonds

Ionic or electrovalent bond
Covalent bond
Coordinate covalent bond
Hydrogen bond

3. Ionic Bond

Ionic bond is formed by transfer of electrons from one atom to another.
It usually forms between a metal and a non-metal.
Metal loses electron to form cation; non-metal gains electron to form anion.
Oppositely charged ions are held by electrostatic attraction.

4. Born-Haber Cycle

Used to explain energetics of ionic compound formation.
Important terms: sublimation energy, ionization energy, dissociation energy, electron affinity, lattice energy.
Formation of ionic compound is favored by low ionization energy, high electron affinity, and high lattice energy.

5. Properties of Ionic Compounds

They are crystalline solids.
They are hard and brittle.
They have high melting and boiling points.
They are generally soluble in water.
They conduct electricity in molten state or in aqueous solution.

6. Covalent Bond

Covalent bond is formed by sharing of electron pairs.
Usually formed between non-metals.
Single bond means one shared pair, double bond means two, triple bond means three.
Octet rule: atoms of second period usually try to complete 8 electrons in outer shell.

7. Lewis Structures

Lewis symbols show valence electrons as dots around atomic symbols.
Useful for writing structures of simple molecules and ions.
Bond pair = shared electron pair.
Lone pair = non-bonding pair of electrons.

8. Coordinate Covalent Bond

Both bonding electrons are donated by one atom.
One species acts as donor and the other as acceptor.
Represented by an arrow from donor to acceptor.
Examples include NH₄⁺ and BF₃ + NH₃.

9. Properties of Covalent Compounds

They usually have low melting and boiling points.
They are poor conductors of electricity.
They are generally insoluble in water.
They dissolve in non-polar solvents like benzene and carbon tetrachloride.

10. Polar Covalent Bond and Dipole Moment

If electronegativity difference is zero or very small, bond is non-polar.
If electronegativity difference exists, bond becomes polar covalent.
Dipole moment measures polarity.
Dipole moment depends on charge and distance of separation.

μ = Q × r

11. Shape and Dipole Moment of Molecules

CO₂ is linear and non-polar overall, so dipole moment is zero.
H₂O is bent and polar.
NH₃ is pyramidal and polar.
BF₃ is planar and symmetric, so net dipole moment is zero.
CCl₄ is tetrahedral and symmetric, so net dipole moment is zero.

12. Fajan's Rules

Ionic bond gains covalent character due to polarization.
Small cation and large anion increase covalent character.
Higher positive charge on cation increases polarization.
More polarization means more covalent character.

13. Bond Parameters

Bond order = number of bonds between two atoms.
Bond length = distance between nuclei of two bonded atoms.
Higher bond order usually means shorter bond length and stronger bond.

14. VSEPR Theory

Electron pairs around central atom arrange themselves to minimize repulsion.
Lone pair-lone pair repulsion is maximum.
Lone pair-bond pair repulsion is less.
Bond pair-bond pair repulsion is minimum.

LP-LP > LP-BP > BP-BP

15. Common Shapes to Memorize

BeCl₂ linear
BF₃ trigonal planar
CH₄ tetrahedral
NH₃ trigonal pyramidal
H₂O bent
SF₆ octahedral

16. Hybridization

Hybridization is mixing of orbitals of comparable energy to form equal number of equivalent hybrid orbitals.
sp gives linear shape.
sp² gives trigonal planar shape.
sp³ gives tetrahedral arrangement.
dsp² gives square planar shape.
dsp³ gives trigonal bipyramidal shape.

17. Sigma and Pi Bonds

Sigma bond is formed by head-on overlap.
Pi bond is formed by sideways overlap.
Single bond = one sigma bond.
Double bond = one sigma + one pi bond.
Triple bond = one sigma + two pi bonds.

18. Resonance

Some molecules cannot be represented by a single Lewis structure.
Such molecules are represented by two or more canonical structures.
Actual structure is resonance hybrid.
Resonance increases stability.
Examples: O₃, CO₃^2-, SO₂, N₂O.

19. Molecular Orbital Theory

Atomic orbitals combine to form molecular orbitals spread over the whole molecule.
Combination gives bonding MO and antibonding MO.
Bonding MO has lower energy; antibonding MO has higher energy.
Electrons fill MOs according to Aufbau principle.

Bond order = (N_b - N_a)/2

20. Important MOT Results

H₂ exists because bond order is 1.
He₂ does not exist because bond order is 0.
O₂ is paramagnetic because it has unpaired electrons.
Molecules with all paired electrons are diamagnetic.

21. NIOS Exam Traps

BF₃ has polar bonds but net dipole moment zero because of symmetry.
Do not confuse coordinate bond formation with ordinary covalent bond formation.
Remember VSEPR shape changes due to lone pairs.
He₂ and Be₂ do not exist on basic MOT treatment because bond order is zero.
O₂ paramagnetism is explained by MOT, not simple Lewis structure.

22. High-Yield Questions to Prepare

Differentiate ionic and covalent bond.
Explain Born-Haber cycle and factors favoring ionic bond formation.
Write Lewis structures of simple molecules and ions.
Predict molecular shapes using VSEPR theory.
Explain hybridization in CH₄, C₂H₄, and C₂H₂.
Write resonance structures and explain resonance hybrid.
Calculate bond order using MOT.

Best exam strategy: remember one example for each bond type, each shape, and each hybridization.