Introduction to Chemical Bonding
Atoms combine to achieve a stable electronic configuration (usually octet or duplet). The force of attraction holding atoms together is called a chemical bond.
Kossel-Lewis Approach (Octet Rule)
Atoms tend to achieve 8 electrons (or 2 for H, He) in their valence shell through electron transfer (ionic) or sharing (covalent).
Ionic Bond
Complete transfer of electrons from one atom to another.
Conditions for Ionic Bond Formation
- Low ionization enthalpy of metal (electropositive).
- High electron gain enthalpy of non-metal (electronegative).
- Large difference in electronegativity (usually > 1.7).
Properties of Ionic Compounds
- High melting and boiling points.
- Soluble in polar solvents (water).
- Conduct electricity in molten or aqueous state.
- Hard and brittle.
- Non-directional bonds.
Covalent Bond
Sharing of electrons between atoms.
Types of Covalent Bonds
- Single bond: One pair of electrons shared (H-H, Cl-Cl).
- Double bond: Two pairs shared (O=O, C=O).
- Triple bond: Three pairs shared (N(triple)N, HC(triple)CH).
Bond Parameters
- Bond length: Distance between nuclei of bonded atoms.
- Bond energy: Energy required to break a bond.
- Bond order: Number of bonds between atoms.
- Bond angle: Angle between two adjacent bonds.
Properties of Covalent Compounds
- Low melting and boiling points.
- Insoluble in water, soluble in organic solvents.
- Do not conduct electricity.
- Directional bonds.
- Can be gases, liquids, or solids.
Coordinate Bond (Dative Bond)
Both shared electrons come from one atom only.
Example: NH_3 -> BF_3 (N donates lone pair to B).
Represented by an arrow from donor to acceptor.
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Electron pairs (bonding and lone pairs) arrange themselves around the central atom to minimise repulsion.
Repulsion Order
Lone pair-Lone pair > Lone pair-Bond pair > Bond pair-Bond pair
Molecular Geometries
| Electron Pairs | Lone Pairs | Shape | Example |
|---|---|---|---|
| 2 | 0 | Linear | BeCl2 |
| 3 | 0 | Trigonal planar | BF3 |
| 3 | 1 | Bent | SO2 |
| 4 | 0 | Tetrahedral | CH4 |
| 4 | 1 | Trigonal pyramidal | NH3 |
| 4 | 2 | Bent/V-shaped | H2O |
| 5 | 0 | Trigonal bipyramidal | PCl5 |
| 6 | 0 | Octahedral | SF6 |
Hybridisation
Mixing of atomic orbitals to form new hybrid orbitals of equivalent energy.
Types of Hybridisation
sp: Linear, bond angle 180. Example: BeCl2, C2H2. sp^2: Trigonal planar, bond angle 120. Example: BF3, C2H4. sp^3: Tetrahedral, bond angle 109.5. Example: CH4, NH3, H2O. sp^3d: Trigonal bipyramidal. Example: PCl5. sp^3d^2: Octahedral. Example: SF6.
Key Points
- Only orbitals of nearly equal energy hybridise.
- Number of hybrid orbitals = number of atomic orbitals mixed.
- Hybrid orbitals overlap to form sigma bonds.
Molecular Orbital Theory (MOT)
Atomic orbitals combine to form molecular orbitals (MO).
Bonding and Antibonding MOs
- Bonding MO: Lower energy (in-phase overlap).
- Antibonding MO: Higher energy (out-of-phase overlap, denoted with asterisk*).
MO Configuration for Homonuclear Diatomic Molecules
Energy order for molecules up to N2: sigma1s, sigma*1s, sigma2s, sigma*2s, sigma2pz, pi2px = pi2py, pi*2px = pi*2py, sigma*2pz.
For O2, F2, Ne2: pi2px = pi2py is lower than sigma2pz.
Bond order: BO = (N_b - N_a)/2
- Positive bond order means stable molecule.
- Higher bond order = shorter and stronger bond.
Magnetic behaviour: Paramagnetic if unpaired electrons, diamagnetic if all paired.
- O2 is paramagnetic (2 unpaired electrons in pi* orbitals).
Hydrogen Bonding
Attraction between hydrogen bonded to a highly electronegative atom (F, O, N) and another electronegative atom.
Types
- Intermolecular: Between different molecules (H2O, HF, NH3). Higher boiling points.
- Intramolecular: Within the same molecule (o-nitrophenol). Lower boiling points.
Effects of Hydrogen Bonding
- Anomalous boiling points of H2O, HF, NH3.
- Ice has less density than water (open cage structure due to H-bonding).
- DNA double helix stabilised by H-bonds.
Worked Examples
Example 1: Predict the shape of NH3 using VSEPR theory. Solution: N has 5 valence electrons, 3 bonded to H and 1 lone pair. Total 4 electron pairs. Shape: trigonal pyramidal (3 bonds + 1 lone pair). Bond angle = 107 degrees.
Example 2: Determine bond order of O2 molecule.
Solution: O2 (16 electrons): MO configuration gives BO = (10-6)/2 = 2. O2 has double bond.
Common Mistakes
- Octet rule exceptions: Incomplete octet (BeCl2, BF3), expanded octet (PCl5, SF6), odd electron species (NO).
- Resonance: Not a real structure between forms; actual structure is a hybrid of all contributing structures.
- VSEPR and lone pairs: Lone pairs exert more repulsion, reducing bond angles from ideal values.
- sigma vs pi bonds: Sigma bonds are stronger, formed by end-on overlap. Pi bonds are weaker, formed by side-on overlap.
ISC Exam Focus
- Theory (70%): Types of bonds, VSEPR theory, hybridisation, MOT, hydrogen bonding.
- Application (30%): Predicting shapes, hybridisation, bond order calculations.
- ISC frequently asks: "Predict the shape of ... using VSEPR" and "Draw MO diagram of ...".
- Comparison of ionic vs covalent compounds.
Self-Test Questions
Q1: Distinguish between ionic and covalent bonds. Answer: Ionic: electron transfer, high MP, soluble in water, conduct electricity. Covalent: electron sharing, low MP, may not dissolve in water, non-conductor.
Q2: Predict the shape and bond angle in H2O using VSEPR theory. Answer: 2 bond pairs + 2 lone pairs = tetrahedral electron pair geometry. Shape = bent/V-shaped. Bond angle = 104.5 degrees (less than 109.5 due to lone pair repulsion).
Q3: What is hybridisation in CH4? Answer: C has sp^3 hybridisation. Four sp^3 hybrid orbitals form sigma bonds with four H atoms. Tetrahedral geometry.
Q4: Calculate bond order of N2.
Answer: N2 (14 electrons): BO = (10-4)/2 = 3. N2 has triple bond.
Q5: Define hydrogen bonding with an example. Answer: Electrostatic attraction between H (bonded to F/O/N) and an electronegative atom. Example: H2O molecules form H-bonds.
Q6: Why is O2 paramagnetic? Answer: O2 MO diagram shows 2 unpaired electrons in pi2px and pi2py orbitals. Unpaired electrons cause paramagnetism.
