Atoms bond together to form more stable arrangements. The type of bond depends on the elements involved: non-metals share electrons (covalent bonding), metals transfer electrons to non-metals (ionic bonding), and metals share a 'sea' of electrons among themselves (metallic bonding).
13.1 Covalent Bonding
Definition
Covalent bond
A covalent bond is a type of chemical bond formed by the sharing of one or more pairs of electrons between two non-metal atoms. Each shared pair constitutes one bond.
Definition
Molecule
A molecule is the smallest unit of a covalently bonded substance that retains the chemical properties of that substance.
Definition
Lewis structure (electron dot structure)
A Lewis structure is a structural formula that shows all valence electrons of every atom in a molecule, including both bonding pairs (shared) and lone pairs (non-bonding).
Non-metal atoms share electrons to achieve a full outer shell. For most atoms this means 8 electrons (the octet rule); for hydrogen and helium it means 2 electrons (duet rule). A single bond = 1 shared pair = 2 electrons. A double bond = 2 shared pairs = 4 electrons. A triple bond = 3 shared pairs = 6 electrons. Electrons not involved in bonding are lone pairs (non-bonding pairs) and are shown as pairs of dots on the atom.
Steps for drawing Lewis structures
- Count the total number of valence electrons from all atoms in the molecule.
- Place the least electronegative atom in the centre (hydrogen is always terminal).
- Connect all atoms to the central atom with single bonds (each uses 2 electrons).
- Distribute remaining electrons as lone pairs, completing the octets of outer atoms first.
- If the central atom still lacks a full octet, convert lone pairs from outer atoms into double or triple bonds with the central atom.
- Check: each atom (except H) should have 8 electrons around it; total electrons used = total valence electrons calculated in step 1.
Worked Example
Draw the Lewis structure for CO₂. (C has 4 valence electrons; each O has 6 valence electrons.)
Given
- C: 4 valence electrons
- O: 6 valence electrons each
- 2 O atoms
Find
Lewis structure for CO₂
Solution
- 1Total valence electrons = 4 + 2(6) = 16 e⁻
- 2C is the central atom (less electronegative than O, and hydrogen is absent).
- 3Place O atoms on each side of C and connect with single bonds: O–C–O (uses 4 e⁻, 12 remaining).
- 4Complete octets on each O atom with 3 lone pairs each: uses 12 e⁻. Now each O has 8 e⁻ but C only has 4.
- 5Convert one lone pair from each O into a double bond with C: O=C=O.
- 6Now C has 8 e⁻ (2+2+2+2 = 4 from double bonds × 2) ✓; each O has 8 e⁻ (4 from double bond + 2 lone pairs × 2) ✓
- 7Total electrons used: 4 (bonds) × 2 + 4 (lone pairs) × 2 = 8 + 8 = 16 ✓
Note
In Grade 10 you need to be able to draw Lewis structures for the following molecules: H₂, Cl₂, HCl, H₂O, NH₃, CH₄, CO₂, N₂. The octet rule applies to all atoms except H (duet) and some exceptions in Period 3 and beyond (not required at Grade 10).
Practice Question
Draw the Lewis structure for NH₃ and identify: (a) the number of bonding pairs, (b) the number of lone pairs on the nitrogen atom.
(4 marks)
13.2 Ionic and Metallic Bonding
Definition
Ionic bond
An ionic bond is a type of chemical bond formed by the electrostatic attraction between oppositely charged ions. Ionic bonds form when a metal atom transfers one or more electrons to a non-metal atom.
Definition
Cation
A cation is a positively charged ion formed when an atom loses one or more electrons. Metals typically form cations. Example: Na → Na⁺ + e⁻.
Definition
Anion
An anion is a negatively charged ion formed when an atom gains one or more electrons. Non-metals typically form anions. Example: Cl + e⁻ → Cl⁻.
Definition
Metallic bonding
Metallic bonding is the electrostatic attraction between the positively charged metal cations (ions in the lattice) and the surrounding sea of delocalised (free) electrons. The delocalised electrons are not associated with any particular atom and can move freely throughout the metal lattice.
Ionic bonds form between metals and non-metals. The metal has 1, 2 or 3 valence electrons which it loses relatively easily (low ionisation energy), becoming a cation with a full outer shell. The non-metal has 5, 6 or 7 valence electrons and gains 1, 2 or 3 electrons to complete its octet, becoming an anion. The opposite charges attract, forming a giant ionic lattice — not individual molecules. The formula of the compound reflects the ratio of ions needed to achieve overall electrical neutrality.
Covalent vs Ionic Bonding
| Property | Covalent Bond | Ionic Bond |
|---|---|---|
| Participants | Non-metal + non-metal | Metal + non-metal |
| Electron behaviour | Electrons shared between atoms | Electrons transferred from metal to non-metal |
| Product formed | Discrete molecules | Giant ionic lattice (no discrete molecules) |
| Electrical conductivity | Generally non-conducting | Conducts when molten or dissolved in water |
Worked Example
Show, using electron dot diagrams, how the ionic compound MgCl₂ is formed from magnesium and chlorine atoms.
Given
- Mg: Group 2 — 2 valence electrons
- Cl: Group 17 — 7 valence electrons
Find
Ionic equation and product
Solution
- 1Mg (2 valence electrons) must lose 2 electrons to achieve a full outer shell (shell 2 with 8 e⁻ from below).
- 2Each Cl atom (7 valence electrons) needs only 1 more electron to complete its octet.
- 3Therefore 1 Mg atom transfers 1 electron to each of 2 Cl atoms:
- 4Mg → Mg²⁺ + 2e⁻
- 52Cl + 2e⁻ → 2Cl⁻
- 6Mg²⁺ + 2Cl⁻ → MgCl₂
- 7The 2+ charge on Mg²⁺ is balanced by two 1− charges on the two Cl⁻ ions — overall charge neutral.
Real World
Table salt (NaCl) has an ionic crystal structure. In solid form, the ions are locked in the lattice and cannot move — so solid NaCl does NOT conduct electricity. When NaCl dissolves in water or is melted, the Na⁺ and Cl⁻ ions become free to move and can carry electric current. This is why salt water (and any dissolved ionic compound) conducts electricity.
Practice Question
Write the formula of the ionic compound formed between Ca²⁺ and O²⁻ ions, and give the name of the compound.
(3 marks)