Atomic structure, the periodic table and chemical bonding

Atomic structure

Every atom consists of a tiny, dense nucleus containing protons (positive charge, relative mass 1) and neutrons (no charge, relative mass 1), surrounded by shells of electrons (negative charge, negligible relative mass ≈ 1/1836).

The atomic number (proton number, Z) is the number of protons. It identifies the element. The mass number A is the total of protons + neutrons. The number of neutrons = A − Z.

In a neutral atom: number of electrons = number of protons.

Isotopes are atoms of the same element (same proton number) with different numbers of neutrons (different mass numbers). E.g. carbon-12 (⁶₁₂C) and carbon-14 (⁶₁₄C) are isotopes of carbon.

Relative atomic mass (Ar) is the weighted mean mass of all naturally occurring isotopes of an element relative to ¹/₁₂ the mass of a carbon-12 atom. Chlorine has Ar ≈ 35.5 because it exists as ~75% ³⁵Cl and ~25% ³⁷Cl.

Electron configuration

Electrons occupy shells (energy levels) around the nucleus. The first shell holds a maximum of 2 electrons; the second and third shells hold up to 8 each. Electrons fill the lowest available shell first.

The number of outer-shell electrons determines chemical behaviour and group membership in the periodic table.

The periodic table

The periodic table arranges elements in order of increasing atomic number. Elements in the same group (vertical column) share the same number of outer-shell electrons and therefore similar chemical properties.

• Group 1 (alkali metals): 1 outer electron; highly reactive; form 1+ ions (e.g. Na⁺, K⁺).
• Group 7 (halogens): 7 outer electrons; form −1 ions or share 1 electron in covalent bonds.
• Group 0 (noble gases): full outer shell; unreactive.

Periods (horizontal rows) represent the number of electron shells. Moving across a period, the atomic number increases by 1.

Metals are on the left and centre; non-metals are on the right. The transition metals in the middle block form coloured compounds and variable oxidation states.

Ionic bonding

Ionic bonding forms between a metal and a non-metal. The metal loses one or more electrons to form a positively charged cation; the non-metal gains electrons to form a negatively charged anion. The electrostatic attraction between oppositely charged ions is the ionic bond.

Dot-and-cross diagrams show electron transfer: draw the outer shells of both atoms, then show electrons moving from metal to non-metal with an arrow (or by placing the electrons in the new ion's outer shell, using crosses for one element and dots for the other).

Example: sodium chloride (NaCl)

• Na (2,8,1) → loses 1e⁻ → Na⁺ (2,8)
• Cl (2,8,7) → gains 1e⁻ → Cl⁻ (2,8,8)

Properties of ionic compounds: high melting and boiling points (strong electrostatic forces between many ions require a lot of energy to overcome); conduct electricity when molten or dissolved in water (ions free to move), but not when solid (ions fixed in lattice).

Covalent bonding

Covalent bonding forms between non-metal atoms. Each atom contributes one or more electrons to form a shared pair (the covalent bond). Each shared pair counts as one bond.

• Single bond: 1 shared pair (e.g. H−H, H−Cl)
• Double bond: 2 shared pairs (e.g. O=O)
• Triple bond: 3 shared pairs (e.g. N≡N)

Dot-and-cross diagrams for covalent molecules (WJEC required): show all outer-shell electrons; bonding pairs are drawn between atoms, lone pairs remain on one atom only.

Simple covalent molecules (e.g. H₂O, CO₂, CH₄): have low melting and boiling points because the intermolecular forces between molecules are weak (not the strong covalent bonds within the molecules). They do not conduct electricity.

Giant covalent structures (e.g. diamond, graphite, silicon dioxide): very high melting points because a huge network of strong covalent bonds must be broken. Diamond does not conduct (no free electrons); graphite does (delocalised electrons between layers).

WJEC required practicals relevant to U1.1

• Flame tests: identify Group 1 and 2 metals by characteristic flame colour (Li = red, Na = yellow/orange, K = lilac, Ca = brick red, Ba = green, Cu = blue-green).
• Electrolysis of copper sulfate (Unit 1 context): demonstrates ionic movement.

Common examiner traps

1. Confusing atomic number and mass number: atomic number = protons; mass number = protons + neutrons.
2. Ionic vs covalent dot-and-cross: in ionic diagrams show square brackets and charge on each ion; in covalent diagrams no brackets/charges.
3. Why ionic compounds don't conduct when solid: ions cannot move — they are in fixed positions in the lattice. Do NOT say "no electrons" — some students confuse this with metallic bonding.
4. Relative atomic mass vs mass number: Ar is a weighted average (often a decimal); mass number is always a whole number.
5. Isotopes have the same number of protons and electrons but different neutron count. They have identical chemical properties but different physical properties (e.g. density, rate of diffusion).