Covalent Bonding

Overview

Covalent bonding occurs when atoms share electrons to achieve stable electron configurations. This typically happens between nonmetal atoms with similar electronegativities.

Types of Covalent Bonds

Single Bond

• 1 shared pair of electrons (2 electrons)
• Represented by one line: —
• Example: H—H, Cl—Cl

Double Bond

• 2 shared pairs of electrons (4 electrons)
• Represented by two lines: =
• Example: O=O, C=O

Triple Bond

• 3 shared pairs of electrons (6 electrons)
• Represented by three lines: ≡
• Example: N≡N, C≡C

Bond Properties

Property	Single	Double	Triple
Bond Order	1	2	3
Bond Length	Longest	Medium	Shortest
Bond Energy	Weakest	Medium	Strongest

Bond Length Comparison

$$\text{C—C (154 pm)} > \text{C=C (134 pm)} > \text{C≡C (120 pm)}$$

Bond Energy Comparison

$$\text{C—C (347 kJ/mol)} < \text{C=C (614 kJ/mol)} < \text{C≡C (839 kJ/mol)}$$

Polar vs Nonpolar Covalent Bonds

Electronegativity Difference ($\Delta EN$)

$\Delta EN$	Bond Type	Electron Distribution
0	Pure covalent	Equal sharing
0.1 - 0.4	Nonpolar covalent	Nearly equal
0.5 - 1.7	Polar covalent	Unequal sharing
> 1.7	Ionic	Electron transfer

Polar Covalent Bond

δ+    δ-
H — Cl

• δ+ = partial positive charge (less electronegative atom)
• δ- = partial negative charge (more electronegative atom)

Dipole Moment ($\mu$)

A measure of bond polarity:

$$\mu = q \times d$$

Where:

• $q$ = magnitude of charge separation
• $d$ = distance between charges
• Unit: Debye (D)

Bond Order

$$\text{Bond Order} = \frac{\text{Bonding electrons} - \text{Antibonding electrons}}{2}$$

For simple molecules:

$$\text{Bond Order} = \text{Number of shared electron pairs}$$

Coordinate (Dative) Bonds

One atom provides both electrons for the shared pair.

Example: Ammonium ion (NH₄⁺)

$$\text{H}_3\text{N:} + \text{H}^+ \rightarrow [\text{NH}_4]^+$$

The lone pair on nitrogen is donated to H⁺.

Resonance

When multiple valid Lewis structures can be drawn for a molecule.

Example: Ozone (O₃)

O=O—O  ↔  O—O=O

• The actual structure is a hybrid
• Each O-O bond is 1.5 order (between single and double)

Example: Carbonate ion (CO₃²⁻)

$$\text{Bond order} = \frac{4}{3} \approx 1.33$$

All three C-O bonds are equivalent.

Sigma ($\sigma$) and Pi ($\pi$) Bonds

Sigma Bonds

• Head-on overlap of orbitals
• Electron density along bond axis
• All single bonds are σ bonds
• Free rotation around the bond

Pi Bonds

• Side-by-side overlap of p orbitals
• Electron density above and below bond axis
• Found in double and triple bonds
• Restricts rotation

Bond Composition

Bond Type	$\sigma$ bonds	$\pi$ bonds
Single (—)	1	0
Double (=)	1	1
Triple (≡)	1	2

Molecular vs Ionic Compounds

Property	Covalent/Molecular	Ionic
Melting Point	Low	High
Boiling Point	Low	High
State at room temp	Often gas/liquid	Solid
Conductivity (solid)	None	None
Conductivity (liquid)	None	Yes
Solubility in water	Varies	Usually soluble

Common Covalent Molecules

Molecule	Formula	Bond Type
Water	H₂O	Polar covalent
Methane	CH₄	Nonpolar covalent
Carbon dioxide	CO₂	Polar bonds, nonpolar molecule
Ammonia	NH₃	Polar covalent
Nitrogen	N₂	Nonpolar covalent (triple bond)