16.3 Chemical Bonding: Ionic, Covalent & Coordinate Bonds, VSEPR, Hybridization & MOT

Chemical Bonding: Ionic, Covalent & Coordinate Bonds, VSEPR, Hybridization & MOT

Chemical bonding explains why atoms combine, what shapes molecules take, and why some substances are polar, magnetic or unusually stable. It links atomic structure (Section 16.2) with intermolecular forces (Section 16.5) and the periodic trends in Section 18.1.

---

1. Why Atoms Bond: The Octet Rule

Atoms bond to reach a lower-energy, more stable arrangement, usually a noble-gas configuration of eight valence electrons (two for hydrogen). The octet rule has exceptions:

• Incomplete octet: $\text{BF}_3$ (boron has 6 valence electrons), $\text{BeCl}_2$ (4).
• Expanded octet: $\text{PCl}_5$ (10), $\text{SF}_6$ (12), possible for Period 3 and heavier elements.
• Odd-electron molecules: $\text{NO}$, $\text{NO}_2$.

---

2. Types of Chemical Bond

Ionic (Electrovalent) Bond

Formed by complete transfer of electrons from a metal of low ionization energy to a non-metal of high electron affinity, followed by electrostatic attraction between the ions. Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$ and $\text{Cl} + e^- \rightarrow \text{Cl}^-$. Ionic compounds have high melting points, conduct when molten or dissolved, and are often soluble in water. Formation is favoured by a large lattice energy.

Covalent Bond

Formed by sharing a pair of electrons: single ($\text{H}-\text{H}$), double ($\text{O}=\text{O}$) or triple ($\text{N} \equiv \text{N}$). A covalent bond between atoms of different electronegativity is polar ($\text{H}^{\delta+}-\text{Cl}^{\delta-}$).

Coordinate Covalent (Dative) Bond

Both electrons of the shared pair come from one atom (the donor) and go to an atom with a vacant orbital (the acceptor). Once formed, it is identical to any other covalent bond.

• $\text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+$ (all four N–H bonds become equivalent)
• $\text{H}_2\text{O} + \text{H}^+ \rightarrow \text{H}_3\text{O}^+$
• $\text{NH}_3 \rightarrow \text{BF}_3$ forms the adduct $\text{H}_3\text{N}\rightarrow\text{BF}_3$

---

3. VSEPR Theory and Molecular Shapes

The valence shell electron pair repulsion theory states that electron pairs around a central atom arrange themselves as far apart as possible. Lone pairs repel more strongly than bonding pairs:

$\text{lone pair–lone pair} > \text{lone pair–bond pair} > \text{bond pair–bond pair}$

The steady fall from $109.5^\circ$ to $107^\circ$ to $104.5^\circ$ as lone pairs replace bond pairs is a classic question. In $\text{CO}_2$, each double bond counts as one electron-pair region.

---

4. Valence Bond Theory: Sigma and Pi Bonds

A covalent bond forms when half-filled atomic orbitals overlap, and the greater the overlap, the stronger the bond.

• Sigma ($\sigma$) bond: head-on overlap along the internuclear axis ($s$–$s$, $s$–$p$, $p$–$p$). Free rotation is possible about a single $\sigma$ bond.
• Pi ($\pi$) bond: sideways overlap of parallel $p$ orbitals above and below the axis. It is weaker than a $\sigma$ bond and prevents rotation, which is why alkenes show cis–trans isomerism.

A single bond is one $\sigma$ bond; a double bond is $1\sigma + 1\pi$; a triple bond is $1\sigma + 2\pi$. Ethyne, $\text{HC}\equiv\text{CH}$, therefore contains $3\sigma$ and $2\pi$ bonds.

---

5. Hybridization

Atomic orbitals of similar energy on the central atom mix to form equivalent hybrid orbitals.

In ethene each carbon is $sp^2$: the hybrid orbitals form $\sigma$ bonds, and the unhybridized $p$ orbitals overlap sideways to form the $\pi$ bond. In ethyne each carbon is $sp$ with two unhybridized $p$ orbitals forming two $\pi$ bonds. Greater $s$-character makes a hybrid orbital shorter and more electronegative, so C–H bonds shorten from ethane to ethene to ethyne.

---

6. Molecular Orbital Theory (MOT)

Atomic orbitals combine to form bonding molecular orbitals (lower energy, electron density between the nuclei) and antibonding orbitals (higher energy, marked $*$). Electrons fill them by the Aufbau principle, Pauli principle and Hund's rule.

$\text{Bond order} = \frac{1}{2}(N_b - N_a)$

where $N_b$ and $N_a$ are the numbers of bonding and antibonding electrons. A higher bond order means a stronger, shorter bond; a bond order of zero means the molecule is not stable.

The paramagnetism of liquid oxygen, which is attracted to a magnet, was a key success of MOT.

---

7. Bond Energy, Bond Length and Dipole Moment

• Bond energy is the energy needed to break one mole of a given bond in the gas phase. It increases with bond order: C–C about 348, C=C about 614 and C≡C about 839 kJ/mol.
• Bond length decreases as bond order increases, and increases with atomic size: H–F < H–Cl < H–Br < H–I.
• Dipole moment $\mu = q \times d$ measures the separation of charge in a polar molecule; its unit is the debye ($1\text{ D} = 3.34 \times 10^{-30}\text{ C}\cdot\text{m}$).

A molecule's dipole moment is the vector sum of its bond dipoles, so shape matters:

A zero dipole moment is evidence of a symmetric shape: $\text{CO}_2$ must be linear, whereas $\text{H}_2\text{O}$, with a large dipole moment, must be bent.