Chemical Bonding

Atoms form bonds because the bonded state has lower energy than the same atoms kept separate. Nature always tends toward lower energy configurations — a bond is simply atoms finding a more stable arrangement together.

The Octet Rule

Most atoms are most stable when surrounded by 8 valence electrons (a full outer shell, matching the configuration of noble gases). Hydrogen and helium are exceptions — they are stable with just 2 electrons (a duet).

• Metals (left side of periodic table) tend to have 1–3 valence electrons and lose them easily
• Nonmetals (right side) tend to have 5–7 valence electrons and gain or share to reach 8
• Noble gases already have 8 valence electrons — they do not need to bond

Ionic bonding occurs when a metal transfers one or more electrons to a nonmetal. The result is two oppositely charged ions held together by electrostatic (Coulombic) attraction.

How It Works

• Metal atom loses electrons → becomes a positively charged cation
• Nonmetal atom gains electrons → becomes a negatively charged anion
• Opposite charges attract, forming an ionic bond
• In solids, ions arrange in a repeating crystal lattice for maximum stability

Key Examples

Compound	Ions formed	e⁻ transfer
NaCl	Na⁺ and Cl⁻	Na loses 1e⁻ to Cl
MgO	Mg²⁺ and O²⁻	Mg loses 2e⁻ to O
CaCl₂	Ca²⁺ and 2×Cl⁻	Ca loses 2e⁻, one to each Cl

Lattice Energy

Lattice energy is the energy released when gaseous ions come together to form a solid ionic lattice. Higher charges and smaller ion sizes produce stronger lattice energy (and higher melting points). MgO has a higher melting point than NaCl because Mg²⁺/O²⁻ have greater charges than Na⁺/Cl⁻.

Properties of Ionic Compounds

• High melting points — strong electrostatic forces hold the lattice together
• Brittle — shifting layers aligns like charges, causing repulsion and cracking
• Conducts electricity when molten or dissolved — ions are free to move
• Does not conduct as a solid — ions are locked in place in the lattice

Covalent bonding occurs between two nonmetals. Neither atom has a strong enough pull to take electrons completely, so they share electron pairs. Each shared pair is simultaneously attracted by both nuclei.

Types of Covalent Bonds

Bond type	Shared pairs	Shared electrons	Bond length	Bond energy
Single bond	1	2 e⁻	Longest	Lowest
Double bond	2	4 e⁻	Medium	Medium
Triple bond	3	6 e⁻	Shortest	Highest

Key Examples

• H₂ — single bond, 2 shared electrons. Each H has a full duet.
• O₂ — double bond, 4 shared electrons. Each O has 8 electrons total.
• N₂ — triple bond, 6 shared electrons. N needs 3 bonds to fill its octet.
• H₂O — 2 single bonds. O shares 1 pair with each H; O has 2 lone pairs.
• CO₂ — 2 double bonds. C shares 2 pairs with each O; written O=C=O.

A Lewis structure (electron dot structure) shows every valence electron in a molecule — both bonding pairs (lines or dots between atoms) and lone pairs (dots on individual atoms).

Step-by-Step Method

Formal Charge

Formal charge helps determine which Lewis structure is best — the structure with formal charges closest to zero is preferred.

VSEPR (Valence Shell Electron Pair Repulsion) theory states that electron groups around a central atom repel each other and arrange themselves to be as far apart as possible, minimizing repulsion.

Geometry Table

Electron groups	Bonding pairs	Lone pairs	Molecular shape	Bond angle	Example
2	2	0	Linear	180°	CO₂, BeCl₂
3	3	0	Trigonal planar	120°	BF₃, SO₃
4	4	0	Tetrahedral	109.5°	CH₄, CCl₄
4	3	1	Trigonal pyramidal	~107°	NH₃, PCl₃
4	2	2	Bent	~104.5°	H₂O, H₂S

Bond Polarity — Electronegativity Difference

When two atoms share electrons, the atom with higher electronegativity (EN) pulls the shared pair closer. If EN values differ significantly, the bond is polar — one end is δ+ and the other is δ−.

ΔEN (electronegativity difference)	Bond type
ΔEN < 0.5	Nonpolar covalent — electrons shared equally
0.5 ≤ ΔEN ≤ 1.7	Polar covalent — unequal sharing, δ+ and δ− ends
ΔEN > 1.7	Ionic — essentially full e⁻ transfer

Molecular Polarity — Shape Matters

A molecule can have polar bonds but still be nonpolar overall if its shape is perfectly symmetrical — the bond dipoles cancel out.

• CO₂ — two polar C=O bonds, linear shape → dipoles point exactly opposite → cancel → nonpolar
• H₂O — two polar O−H bonds, bent shape → dipoles do NOT cancel → polar molecule
• CCl₄ — four polar C−Cl bonds, tetrahedral shape → perfectly symmetric → nonpolar
• NH₃ — three polar N−H bonds, trigonal pyramidal → asymmetric → polar molecule

Metallic bonding occurs between metal atoms. Metal atoms release their valence electrons into a shared "sea" of delocalized electrons that moves freely throughout the entire solid. The positive metal cations are embedded in and held together by this electron sea.

Properties Explained by Metallic Bonding

• Electrical conductivity — delocalized electrons flow freely under a voltage
• Thermal conductivity — mobile electrons transfer kinetic energy quickly
• Malleability and ductility — layers of cations can slide past each other; the electron sea readjusts, preventing fracture
• Metallic luster — free electrons absorb and re-emit light across the visible spectrum

Mistake	What to do instead
Confusing bond polarity with molecular polarity	Check the shape first. Symmetric molecules cancel out their polar bonds — CO₂ has polar bonds but is nonpolar overall.
Ignoring lone pairs in VSEPR	Lone pairs count as electron groups. NH₃ has 4 electron groups (3 bonds + 1 LP) → trigonal pyramidal, not trigonal planar.
Counting a double or triple bond as multiple groups	A multiple bond counts as ONE electron group in VSEPR. CO₂ has 2 groups (two double bonds) → linear.
Forgetting to account for charge in Lewis structures	Add 1 electron for each negative charge; remove 1 for each positive charge before drawing.
Using formal charge = 0 as the only criterion	The best structure minimises formal charges AND places negative formal charges on the most electronegative atoms.
Saying "ionic compounds conduct electricity"	They only conduct when molten or dissolved. Solid ionic compounds do not conduct because ions are fixed in the lattice.