Inside the atom

An atom has a tiny central nucleus holding protons and neutrons, surrounded by electrons in shells. The nucleus is minute compared with the whole atom, yet it holds almost all of its mass.

Watch out: shells fill 2, then 8, then 8 for the first 20 elements. The number of electrons in the outer shell equals the group number (for groups 1–7).

Charges, masses & the two numbers

Each particle has a relative charge and relative mass:

Two numbers describe any atom:

• Atomic number (Z) = number of protons (the bottom number; it identifies the element).
• Mass number (A) = number of protons + neutrons (the top number).

Your turn — counting particles

How the model changed

The model of the atom changed as new evidence — the discovery of subatomic particles — came in:

Isotopes and A_r

Isotopes are atoms of the same element (same protons) with different numbers of neutrons — so the same atomic number but different mass numbers.

Because elements are mixtures of isotopes, the relative atomic mass (A_r) is a weighted mean — which is why many A_r values are not whole numbers.

Common slip: A_r is a weighted average, not a simple (35+37)/2 = 36. The more abundant isotope pulls the mean towards it.

Your turn — relative atomic mass

Reading the periodic table

The modern table arranges elements in order of increasing atomic number. Horizontal rows are periods; vertical columns of similar elements are groups.

Electronic config & position: the group number gives the number of outer-shell electrons; the period number gives the number of occupied shells. So Na (2,8,1) is in group 1, period 3.

How Mendeleev built his table

• Mendeleev arranged the known elements using the properties of the elements and their compounds.
• He left gaps and used the pattern to predict the existence and properties of elements not yet discovered (e.g. "eka‑silicon", later germanium).
• He thought he had ordered them by increasing relative atomic mass — but this wasn't always true, because the relative abundance of isotopes means some pairs come in the "wrong" mass order.

The modern table fixes this by ordering on atomic number (proton number) instead of atomic mass.

Config & position

Ionic bonding: transfer of electrons

An ionic bond forms when electrons transfer from a metal atom to a non‑metal atom, making charged ions: positive cations and negative anions. An ion is an atom (or group of atoms) with a charge.

Ions form to reach a full outer shell. For groups 1, 2, 6 and 7 you can predict the charge from the group:

• Group 1 → 1+ · Group 2 → 2+ (lose electrons)
• Group 6 → 2− · Group 7 → 1− (gain electrons)

Naming & formulae: ‑ide = a simple ion of one element (chloride Cl⁻, oxide O²⁻); ‑ate = an ion containing oxygen (sulfate SO₄²⁻, nitrate NO₃⁻, carbonate CO₃²⁻). Balance the charges to deduce the formula, e.g. Mg²⁺ + Cl⁻ → MgCl₂.

Giant ionic lattices

Ionic compounds don't exist as little molecules — they build a giant lattice: a regular 3‑D arrangement of ions held by strong electrostatic forces (ionic bonds) acting in all directions between oppositely‑charged ions.

Why it matters: because so many strong bonds must be broken, ionic compounds have high melting and boiling points. They conduct only when molten or dissolved (ions free to move) — never as a solid.

Predict the bonding

Covalent bonding: sharing electrons

A covalent bond is a shared pair of electrons between two non‑metal atoms. Covalent bonding makes molecules. Atoms and small molecules are tiny — of the order of 10⁻¹⁰ m across.

You should be able to draw dot‑and‑cross diagrams for: H₂, HCl, H₂O, CH₄, O₂ and CO₂ (O₂ and CO₂ use double bonds — two shared pairs).

Simple molecular properties: the molecules are held to each other by weak intermolecular forces (not the strong covalent bonds inside them), so these substances have low melting and boiling points and do not conduct electricity.

Giant covalent: diamond, graphite & more

Diamond and graphite are both forms of carbon and both are giant covalent (macromolecular) structures — millions of atoms joined by strong covalent bonds.

• Diamond: each carbon bonds to 4 others in a rigid 3‑D network → very hard, very high melting point → used in cutting tools. No free electrons → does not conduct.
• Graphite: each carbon bonds to 3 others in flat layers; the spare (4th) electron per atom is delocalised. Layers slide → used as a lubricant; delocalised electrons → conducts, so used for electrodes.
• Graphene (one layer of graphite) and fullerenes (e.g. C₆₀) — strong, conducting forms of carbon with their own structures and uses.
• Polymers (e.g. poly(ethene)) are large molecules with long chains of carbon atoms.

Why graphite conducts

Metallic bonding

In a metal, the atoms pack into a giant lattice of positive ions sitting in a "sea" of delocalised electrons (the outer electrons that have left the atoms). The strong attraction between the positive ions and this sea is the metallic bond.

Properties explained: the delocalised electrons can carry charge → metals conduct electricity (and heat). The layers of ions can slide over each other without breaking the bonding → metals are malleable. The strong bonds give most metals high melting points.

Name the bonding

Tap the type of bonding that holds each substance together.

Structure decides properties

Edexcel wants you to classify a substance as ionic, simple molecular (covalent), giant covalent or metallic, and explain its properties from its structure and bonding:

Models have limits: dot‑and‑cross, ball‑and‑stick and 2‑D/3‑D drawings each show some things well but mislead on others (e.g. dot‑and‑cross shows bonding but not the true 3‑D shape or that electrons are identical).

Explain the property

Relative formula mass (M_r)

The relative formula mass (M_r) of a compound is found by adding up the A_r values of every atom in its formula.

Also in Topic 1: the law of conservation of mass, finding masses of reactants/products from a balanced equation, and concentration in g dm⁻³.

Your turn — relative formula mass

Moles & empirical formulae HT

These calculation skills in Topic 1 are Higher‑tier only:

• One mole = the Avogadro number of particles, 6.02 × 10²³, with a mass equal to the relative particle mass in grams.
• Convert between mass, moles and number of particles using moles = mass ÷ M_r.
• Work out an empirical formula from reacting masses or percentage composition, and link empirical ↔ molecular formula.
• Deduce the stoichiometry of a reaction from masses, and explain why the limiting reactant (not in excess) controls the mass of product.

Your turn — moles HT

Conductor or not?

Tap a substance, then tap the box for whether it conducts electricity as a solid.

The key ideas to know

You've covered all of Edexcel Topic 1 — atomic structure, the periodic table, ionic/covalent/metallic bonding, types of substance and calculations involving masses. Press Finish to see your score.