Atoms rarely exist in isolation. They combine through chemical bonds to form molecules and compounds. The nature of bonding determines a substance's structure, shape, and properties. This chapter covers ionic, covalent, and coordinate bonds; molecular geometry; and the theories that explain bonding.
---
Key Concepts
Why Do Atoms Bond?
Atoms bond to achieve a lower energy state. This is typically associated with acquiring a stable octet configuration (8 electrons in the valence shell) — the Octet Rule. Exceptions include hydrogen (duet), and certain molecules like BF3 (6), PCl5 (10), and SF6 (12) which expand beyond an octet.
Types of Chemical Bonds
Ionic (Electrovalent) Bond
Formed by the complete transfer of one or more electrons from a metal to a non-metal. The resulting oppositely charged ions (cations and anions) attract each other electrostatically. Example: Na + Cl → Na+ + Cl- → NaCl. Ionic compounds have high melting points, are soluble in polar solvents, and conduct electricity in solution or melt.
- Covalent Bond
- Formed by the sharing of electron pairs between two atoms (usually non-metals). A shared pair is a bond pair; unshared pairs are lone pairs.
- Single bond: 1 shared pair; Double bond: 2 shared pairs; Triple bond: 3 shared pairs.
- Bond order = number of bonds between two atoms. Higher bond order = shorter, stronger bond.
Coordinate (Dative) Bond
A type of covalent bond where both electrons in the shared pair come from one atom (the donor). E.g., in NH4+, N donates both electrons to H+.
- 1.Lewis Structures
- 2.Lewis dot structures show the arrangement of bonding and lone pairs. Steps:
- 3.Count total valence electrons.
- 4.Arrange atoms (least electronegative atom is usually central).
- 5.Distribute electrons to satisfy octets (start from outer atoms).
- 6.Form double/triple bonds if needed.
Formal Charge
Formal charge = (Valence electrons of atom) - (Non-bonding electrons) - 1/2(Bonding electrons). The best Lewis structure minimises formal charges.
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Electron pairs (bonding and lone) around a central atom arrange themselves to minimise repulsion. Repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
- Key shapes:
- 2 bond pairs, 0 lone pairs: Linear (BeCl2, CO2) — 180 degree
- 3 bp, 0 lp: Trigonal planar (BF3) — 120 degree
- 4 bp, 0 lp: Tetrahedral (CH4) — 109.5 degree
- 3 bp, 1 lp: Trigonal pyramidal (NH3) — ~107 degree
- 2 bp, 2 lp: Bent/V-shaped (H2O) — ~104.5 degree
- 5 bp, 0 lp: Trigonal bipyramidal (PCl5)
- 6 bp, 0 lp: Octahedral (SF6)
- Valence Bond Theory (VBT)
- A covalent bond forms when two half-filled orbitals of two atoms overlap. Greater overlap = stronger bond.
- Sigma (sigma) bond: End-to-end (axial) overlap. Rotation is possible.
- Pi (pi) bond: Sideways (lateral) overlap of p orbitals. Prevents rotation. Pi bonds are weaker than sigma bonds.
- Hybridisation
- Mixing of atomic orbitals to form new hybrid orbitals of equivalent energy and shape:
- sp: 2 orbitals, linear, 180 degree (e.g., BeCl2, C2H2)
- sp2: 3 orbitals, trigonal planar, 120 degree (e.g., BF3, C2H4)
- sp3: 4 orbitals, tetrahedral, 109.5 degree (e.g., CH4, NH3, H2O)
- sp3d: Trigonal bipyramidal (e.g., PCl5)
- sp3d2: Octahedral (e.g., SF6)
Molecular Orbital Theory (MOT) — Basics
Atomic orbitals combine to form molecular orbitals (MOs). Bonding MOs (sigma, pi) are lower in energy than the original AOs; antibonding MOs (sigma · , pi · ) are higher. Bond order = (bonding electrons - antibonding electrons) / 2. If bond order > 0, molecule is stable.
Polarity of Bonds and Molecules
A bond is polar if the two atoms have different electronegativities. A molecule is polar (has a dipole moment) if the vector sum of bond dipoles is non-zero. CO2 is non-polar (linear, dipoles cancel); H2O is polar (bent, dipoles add).
Hydrogen Bond
An intermolecular (or intramolecular) attraction between a H atom covalently bonded to a highly electronegative atom (F, O, N) and another electronegative atom with a lone pair. Explains the anomalously high boiling point of H2O.
---
Worked Examples
Draw the Lewis structure of CO2.
C has 4 valence electrons; each O has 6. Total = 4 + 2(6) = 16 e. Arrangement: O=C=O. Each O has 2 lone pairs, C has 0. Structure: :O=C=O: (16 electrons used). Shape: Linear.
Predict the shape of NH3 using VSEPR.
N has 4 electron pairs (3 bp + 1 lp). Electron geometry: tetrahedral. The lone pair compresses the bond angle from 109.5 degree to ~107 degree. Molecular shape: trigonal pyramidal.
What is the hybridisation of carbon in ethylene (C2H4)?
Each C forms a double bond (one sigma + one pi) with the other C and two sigma bonds with H. Total 3 sigma bonds around each C → sp2 hybridisation. Bond angles ~120 degree.
Calculate the bond order of O2 using MOT.
O2 has 16 electrons. MO configuration: (sigma 1s)2 (sigma · 1s)2 (sigma 2s)2 (sigma · 2s)2 (sigma 2p)2 (pi 2p)4 (pi · 2p)2
Bonding electrons = 10, Antibonding = 6. Bond order = (10-6)/2 = 2. (O2 also has 2 unpaired electrons → paramagnetic.)
Why is the bond angle in H2O (104.5 degree) less than in NH3 (107 degree)?
H2O has 2 lone pairs on O; NH3 has 1 lone pair on N. Greater lone pair repulsion in H2O compresses the H-O-H angle more than in NH3.
Identify the formal charge on each atom in the nitrate ion (NO3-).
In one resonance structure with 1 N=O double bond and 2 N-O single bonds: N has formal charge = 5 - 0 - 6 = -1... (using the resonance hybrid, average formal charge on O = -2/3 and N = 0 for the hybrid). The formal charge approach helps identify the best resonance contributor.
Which molecule has the highest dipole moment: BF3, NF3, or NH3?
BF3 is trigonal planar — dipoles cancel, dipole moment = 0. NF3 and NH3 are both pyramidal, but in NH3 the N-H bond dipoles and lone pair both point in the same direction, while in NF3 the N-F dipoles partially cancel the lone pair effect. NH3 has the highest dipole moment among the three.
---
Common mistakes
Students confuse electron geometry (based on all electron pairs) with molecular shape (based on atoms only). For example, H2O has tetrahedral electron geometry but a bent molecular shape. Also, remember that in hybridisation, lone pairs do occupy hybrid orbitals — so NH3 is sp3, not sp2.
---
Summary
Chemical bonding ranges from the purely electrostatic (ionic) to the shared (covalent). VSEPR predicts shape from electron pair repulsion. Hybridisation (VBT) explains observed bond angles. MOT explains properties like paramagnetism and bond order. The polarity and hydrogen bonding of molecules ultimately dictate their physical properties.