Atomic Structure and the Periodic Table

Almost everything else in GCSE Chemistry depends on what you learn here. How an atom is built, and in particular how its electrons are arranged, is what decides how an element behaves and where it sits on the periodic table. Get that straight now and the rest of the course becomes much easier to follow.

1Atoms, Elements & Compounds

Chemical Symbols and Formulae

Every element has a unique symbol — one or two letters, where the first letter is always a capital and any second letter is always lowercase.

A chemical formula shows how many atoms of each element are in one unit of a substance. Subscript numbers (written below the line) indicate more than one atom of that element:

2Chemical Equations & Conservation of Mass

Word Equations

A word equation describes a chemical reaction using the names of reactants (left) and products (right), separated by an arrow:

Magnesium + Oxygen → Magnesium oxide

Naming rule: when a metal reacts with a non-metal, the metal's name is unchanged and the non-metal's name changes to end in -ide.

Balanced Symbol Equations

Atoms are never created or destroyed in a chemical reaction — they are rearranged. This means the number of atoms of each element must be the same on both sides of the equation. We balance equations by adding coefficients (large numbers in front of formulae) — never by changing the subscripts inside a formula.

Conservation of Mass

Because atoms are neither created nor destroyed, the total mass of the products always equals the total mass of the reactants. This is the Law of Conservation of Mass.

Half Equations and Ionic Equations H

At Higher Tier you may also need to write half equations and ionic equations.

A half equation shows either the oxidation or reduction step in a reaction, including electrons (e⁻). Example — sodium losing an electron:

e.g. Na → Na⁺ + e⁻

An ionic equation shows only the ions and species that actually change — spectator ions (present but unchanged) are cancelled out. Example — the ionic equation for neutralisation:

e.g. H⁺(aq) + OH⁻(aq) → H₂O

You will meet half equations and ionic equations throughout the course wherever they are relevant (e.g. electrolysis, displacement reactions, neutralisation).

3Separating Mixtures

A mixture contains two or more substances that are not chemically joined, so each keeps its own properties. That means we can pull them apart using physical processes — by particle size, solubility, or boiling point — with no chemical reaction and no new substances made. You need to be able to describe each method below, and choose the right one for a given mixture.

Several of these methods depend on whether a substance dissolves, so it is worth being clear on a few closely-related words before we start.

Filtration

Filtration separates an insoluble solid from a liquid — for example sand from water. The mixture is poured into a filter funnel lined with filter paper. The liquid passes through the tiny holes in the paper and is collected as the filtrate, while the solid particles are too big to pass through and stay on the paper as the residue.

Crystallisation

Crystallisation separates a soluble solid that is dissolved in a solution — for example salt from salt water. The solution is gently heated in an evaporating basin — usually resting on a beaker of hot water as a water bath, which gives slow, even heat so the crystals do not spit or break down — and the water evaporates away. The dissolved solid cannot evaporate, so as the solution becomes more concentrated it is left behind, forming crystals in the basin.

Simple distillation

Simple distillation separates the liquid (solvent) from a solution — for example pure water from salt water. The solution is heated until the liquid boils. Its vapour travels into a water-cooled condenser, where it cools back into a liquid and drips out as the pure distillate; the dissolved solid is left behind in the flask. It only works cleanly when the parts of the mixture have very different boiling points.

Fractional distillation

Fractional distillation separates a mixture of two or more liquids that mix completely but have different boiling points — for example water from ethanol.

Chromatography

Chromatography separates substances that are dissolved in a solvent, such as the coloured dyes in an ink.

4History of the Atomic Model

Our model of the atom has changed over time as new experimental evidence emerged. Each new model was developed because the previous one could not explain new observations. This is an important example of how science works.

• Early 1800s

  John Dalton (name not required by spec)

  Before the discovery of the electron, atoms were thought to be tiny spheres that could not be divided.

• 1897

  

  Thomson's Plum Pudding Model

  J.J. Thomson name required

  Discovered the electron — a negatively charged particle inside the atom. Proposed the plum pudding model: the atom is a ball of positive charge with negative electrons embedded in it.

• 1911

  Rutherford's nuclear model

  Ernest Rutherford name required

  Fired positively charged alpha particles at a thin gold foil. Expected all particles to deflect slightly (as the Plum Pudding model predicted). Instead:

  • Most passed straight through → the atom is mostly empty space.
  • A tiny proportion deflected at large angles → there is a tiny, dense, positively charged nucleus at the centre.

  This disproved the Plum Pudding model and led to the nuclear model of the atom.

  Straight through — most particles Deflected — a few Bounced back — very rare

  Press play to fire a beam of identical alpha particles at the gold foil.

  ▶ Play animation
• 1913

  Niels Bohr name required

  Adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances (fixed energy levels/shells). The theoretical calculations of Bohr agreed with experimental observations.

• ~1920

  Discovery of the Proton (no specific name required)

  Later experiments showed that the positive charge of the nucleus could be subdivided into a whole number of smaller particles, each with the same amount of positive charge. These particles were named protons.

• 1932

  James Chadwick name required

  Provided experimental evidence for the existence of neutrons within the nucleus — particles with the same mass as a proton but no electrical charge. This was about 20 years after the nucleus became an accepted scientific idea.

5Structure of the Atom

A carbon atom: 6 protons and 6 neutrons in the nucleus, with 6 electrons arranged across two shells (2, 4). Not to scale.

Subatomic Particles

Why Atoms are Electrically Neutral

In a neutral atom, the number of electrons always equals the number of protons. Each proton has charge +1 and each electron has charge −1 — these cancel out exactly, giving a net charge of zero.

Atomic Number and Mass Number

On the Periodic Table, for sodium: mass number = 23, atomic number = 11. So: protons = 11  |  neutrons = 23 − 11 = 12  |  electrons = 11

Practice — Protons, Neutrons and Electrons

Fill in the Protons, Neutrons and Electrons columns, then click Check answers.

6Ions

Working Out Particles in an Ion

The atomic number still gives the number of protons (protons are never gained or lost). The ion's charge tells you how many electrons have been gained or lost.

7Size of Atoms & the Nucleus

This means the atom is mostly empty space — a key conclusion from Rutherford's experiment. Despite being tiny, the nucleus accounts for virtually all of the atom's mass.

8Isotopes & Relative Atomic Mass

Chlorine Isotopes — A Comparison

Relative Atomic Mass (A_r)

Because an element often exists as a mixture of isotopes, the A_r shown on the Periodic Table is a weighted average mass of all its naturally occurring isotopes, taking into account their relative abundances. This is why chlorine has an A_r of 35.5 — it sits between its two isotopes (Cl-35 and Cl-37).

Calculating A_r from Isotope Abundances H

Higher Tier students must be able to calculate A_r when given the masses and percentage abundances of each isotope:

A_r = sum of (isotope mass × % abundance)100

9Electronic Structures

Electrons occupy energy levels (shells) surrounding the nucleus. Shells fill from the innermost (lowest energy) outwards. The arrangement of electrons is called the electronic structure or electron configuration.

First 20 Elements

Hover over any element to see its electron shell diagram.

Metals are generally on the left and centre of the periodic table. Non-metals are generally on the right. Hydrogen is the main exception: it is placed on the left but behaves as a non-metal.

Link to the Periodic Table

10Development of the Periodic Table

Before scientists understood atomic structure, they tried to classify known elements by atomic mass (atomic weight). The table evolved through several key stages as new evidence emerged — a key example of how science works.

• Before 1860s

  Early Classification Attempts (no specific scientist names required)

  Before the discovery of protons, neutrons and electrons, scientists attempted to classify the elements by arranging them in order of their atomic weights. The early periodic tables were incomplete — not all elements had been discovered — and some elements were placed in inappropriate groups if the strict order of atomic weights was followed.

• 1864

  John Newlands — Law of Octaves (name not required by spec)

  Newlands arranged the 56 known elements in order of atomic weight and noticed that every eighth element had similar properties — he called this the Law of Octaves. His table was criticised because:

  • The pattern broke down for heavier elements — some groups contained elements with very different properties
  • He mixed up metals and non-metals in the same groups
  • He left no gaps for undiscovered elements, so newly found elements did not fit
• 1869

  Dmitri Mendeleev name required

  Mendeleev overcame some of the problems with earlier tables by making two key changes:

  • He left gaps for elements he thought had not yet been discovered
  • In some places he changed the order based on atomic weights so that elements with similar chemical properties stayed in the same group

  Mendeleev was able to predict the properties of undiscovered elements from patterns in his table. Elements with the properties he had predicted were subsequently discovered and filled the gaps (e.g. gallium, germanium, scandium) — providing strong evidence for his arrangement, and leading to its acceptance.

  💡 Exam Tip — Why Mendeleev's Table Was Accepted

  Make sure you say both parts: the elements fitted the gaps he left, and their properties matched his predictions.

• Late 1800s – early 1900s

  Discovery of Subatomic Particles (see atomic model timeline — names required there)

  The discovery of electrons, protons and neutrons (covered in Section 5) revealed why elements in the same group behave similarly — they have the same number of outer shell electrons. This gave the periodic table a deeper theoretical basis.

• Modern

  The Modern Periodic Table — Ordered by Atomic Number (no specific scientist name required)

  In the modern periodic table, elements are arranged in order of atomic number (number of protons), not atomic mass. This resolves the occasional inconsistencies in Mendeleev's arrangement.

  ✅ The Role of Isotopes

  Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons — they have different masses but belong in the same position on the table. Because of isotopes, average atomic mass does not always increase in the same order as atomic number.

  Example: argon (A_r = 39.9, atomic number 18) comes before potassium (A_r = 39.1, atomic number 19). Strict atomic mass ordering would swap them — placing potassium in the wrong group. Ordering by atomic number resolves this.

Groups and Periods in the Modern Table

Metals and Non-metals

Elements that react to form positive ions are metals. Elements that do not form positive ions are non-metals. The majority of elements are metals. Metals are found to the left and towards the bottom of the periodic table; non-metals are found towards the right and top.

The position of an element in the periodic table is linked to its atomic structure. Metals have few outer-shell electrons (typically 1–3) and tend to lose them, forming positive ions. Non-metals have more outer-shell electrons (typically 4–7) and tend to gain or share electrons.

Because reactivity depends on how readily an element gains or loses electrons, the reactions of elements are directly related to the arrangement of electrons in their atoms — and therefore to their atomic number and position in the table.

11Group 1 — The Alkali Metals

The alkali metals are the elements in Group 1: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). At GCSE we focus mainly on Li, Na and K.

Physical Properties

Reactions with Water

Alkali metals react with water to produce a metal hydroxide and hydrogen gas. The hydroxide dissolves to give an alkaline solution (hence the group name).

alkali metal + water → metal hydroxide + hydrogen

Reactions with Oxygen

Alkali metals react with oxygen in the air to form metal oxides. This is why they tarnish rapidly when freshly cut — the shiny surface turns dull as the oxide layer forms.

The oxides formed are ionic compounds (metal oxide ion, O²⁻, combined with the metal ion). All three metals are stored under oil to prevent reaction with oxygen and moisture in the air.

Reactions with Chlorine

Alkali metals react vigorously with chlorine gas to form metal chlorides — white, ionic solids that dissolve in water to give a neutral solution:

Ionic Compounds with Non-metals

When alkali metals react with non-metals they form ionic compounds. The alkali metal loses its one outer electron to achieve a full outer shell, forming an ion with a charge of +1. These compounds are:

• Always white solids
• Soluble in water — they dissolve to give colourless solutions
• Held together by electrostatic attraction between oppositely charged ions

Example: sodium reacting with chlorine to form sodium chloride:

2Na(s) + Cl₂(g) → 2NaCl(s)

Reactivity Trend Down Group 1

Reactivity increases going down Group 1: Li → Na → K → Rb → Cs

12Group 7 — The Halogens

The halogens (Group 7) are non-metals: fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). They all have 7 electrons in their outer shell and all exist as diatomic molecules — pairs of atoms bonded together (F₂, Cl₂, Br₂, I₂).

Physical Properties at Room Temperature

Going down Group 7, melting and boiling points increase. The relative molecular mass of the diatomic molecules increases down the group — larger molecules have stronger intermolecular forces, requiring more energy to separate.

📌 Intermolecular forces — the weak forces between molecules that have to be overcome to melt or boil a simple molecular substance — are covered in detail in C2 — Bonding, Structure and the Properties of Matter.

Reactions with Metals — Ionic Compounds

Halogens react with metals to form ionic compounds. Each halogen atom has 7 outer electrons and gains one electron to complete its outer shell, forming a halide ion with a charge of −1.

Example: sodium reacting with chlorine to form sodium chloride (ionic compound):

2Na(s) + Cl₂(g) → 2NaCl(s)

Reactions with Non-metals — Molecular Compounds

When halogens react with non-metals, atoms share electrons, forming molecular (covalent) compounds. Example — hydrogen reacting with chlorine to form hydrogen chloride:

H₂(g) + Cl₂(g) → 2HCl(g)

HCl dissolves in water to form hydrochloric acid.

Reactivity Trend Down Group 7

Reactivity decreases going down Group 7: F₂ > Cl₂ > Br₂ > I₂

Displacement Reactions

A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.

Now try it yourself — pick any halogen and any halide solution to see whether a reaction happens, what you would observe, and the equations.

13The Transition Elements T

The transition elements (transition metals) occupy the central block of the periodic table, between Groups 2 and 3. They are all metals with several distinctive properties that set them apart from the Group 1 and Group 2 metals.

Properties of Transition Metals

AQA names six transition metals to illustrate these properties: chromium (Cr), manganese (Mn), iron (Fe), cobalt (Co), nickel (Ni) and copper (Cu) — you should be able to use these as your examples.

Transition Metals vs Group 1 Alkali Metals

The transition metals are very different from the Group 1 alkali metals. The spec requires you to be able to describe and explain these differences:

Variable Ion Charges

Unlike Group 1 metals (always +1) and Group 2 metals (always +2), transition metals can form ions with different positive charges. The charge is shown by a Roman numeral:

Transition Metals as Catalysts

Uses of Key Transition Metals (beyond the spec — useful context, not required knowledge)

14Noble Gases (Group 0)

The noble gases occupy Group 0 — the rightmost column of the periodic table. They are all colourless, odourless gases at room temperature.

Why Noble Gases Are Unreactive

Noble gases are chemically inert (very unreactive). Each noble gas atom already has a full outer electron shell, so it has no tendency to gain, lose, or share electrons. There is no chemical "drive" to react.

Trend in Boiling Points

The boiling points of the noble gases increase with increasing relative atomic mass going down Group 0. This is because larger atoms have stronger intermolecular forces, requiring more energy to overcome.

Uses of Noble Gases

Their inertness makes noble gases valuable where reactions with surrounding materials must be avoided: