Subatomic particles, isotopes, electronic configuration and the history of the atomic model — everything you need for GCSE Chemistry.
Atomic structure is the foundation of the entire GCSE Chemistry course. Without understanding atoms — what they contain, how they're arranged and why they behave as they do — bonding, reactivity and the periodic table all become much harder. This guide builds that understanding from scratch, covering subatomic particles, the development of the atomic model, isotopes, electronic configuration and relative atomic mass.
Every atom consists of a tiny, dense nucleus surrounded by electrons. The nucleus contains protons and neutrons. The electrons occupy shells (energy levels) around the nucleus.
In a neutral atom: protons = electrons. If an atom gains electrons it becomes a negative ion. If it loses electrons it becomes a positive ion. The number of protons never changes in a chemical reaction — only electrons are transferred or shared.
Each element on the periodic table shows two numbers. The atomic number (proton number) is the smaller number — it gives the number of protons. The mass number (nucleon number) is the larger number — it gives the total number of protons plus neutrons.
So sodium has 11 protons, 11 electrons and 12 neutrons.
Isotopes are atoms of the same element with different numbers of neutrons. They have the same atomic number (same element, same number of protons) but different mass numbers.
For example, carbon has three naturally occurring isotopes: carbon-12 (6 protons, 6 neutrons), carbon-13 (6 protons, 7 neutrons) and carbon-14 (6 protons, 8 neutrons). All three are carbon — they have 6 protons and 6 electrons — but they differ in the number of neutrons in the nucleus.
Isotopes of the same element have identical chemical properties because chemical behaviour depends on the number and arrangement of electrons — which is the same for all isotopes. They differ only in physical properties such as mass and density, and in radioactivity (unstable isotopes are radioactive).
Because most elements exist as mixtures of isotopes, the relative atomic mass (Ar) is a weighted average — it takes into account both the mass and the relative abundance of each isotope.
Example: Chlorine exists as 75% chlorine-35 and 25% chlorine-37. Ar = (35 × 75 + 37 × 25) ÷ 100 = (2625 + 925) ÷ 100 = 3550 ÷ 100 = 35.5. This is why chlorine's Ar on the periodic table is 35.5, not a whole number.
Electrons occupy shells at different energy levels around the nucleus. The first shell holds a maximum of 2 electrons. The second shell holds a maximum of 8. The third shell also holds 8 at GCSE level (in reality it can hold more, but GCSE only covers elements up to calcium where this simplification holds).
Electrons fill the lowest energy shell first. Once a shell is full, electrons go into the next shell out.
The number of outer electrons equals the group number in the periodic table. Sodium has 1 outer electron → Group 1. Chlorine has 7 outer electrons → Group 7. The number of shells equals the period number. Sodium has 3 shells → Period 3.
You can work out electronic configuration from position in the periodic table without memorising it. Period = number of shells. Group = number of outer electrons. Sodium is in Period 3, Group 1 → 3 shells, 1 outer electron → 2,8,1. Chlorine is in Period 3, Group 7 → 3 shells, 7 outer electrons → 2,8,7. This works for all elements in Groups 1–7 up to Period 4.
Our understanding of atomic structure has changed significantly over time as new experimental evidence has emerged. GCSE Chemistry requires you to know the key stages of this development and the evidence behind each change.
John Dalton proposed that matter is made of tiny, indivisible spheres — like tiny solid balls. He had no knowledge of subatomic particles. This model explained why elements combine in fixed ratios but could not explain electrical phenomena or emission spectra.
J.J. Thomson discovered the electron in 1897 through cathode ray experiments — he showed that atoms contain negatively charged particles much lighter than atoms themselves. His model proposed atoms as a sphere of positive charge with electrons embedded throughout — like plums in a pudding. This was the first model to include subatomic particles.
Ernest Rutherford performed the famous gold foil experiment. He fired positively charged alpha particles at a very thin sheet of gold foil and observed where they went. Most passed straight through (expected if Thomson's model were correct). But a small fraction were deflected at large angles, and a tiny fraction bounced straight back. This was completely unexpected — Rutherford famously said it was "as if you fired artillery shells at tissue paper and they came back and hit you."
The only explanation: atoms are mostly empty space, with a tiny, dense, positively charged nucleus at the centre. Most alpha particles miss the nucleus entirely and pass through. Those that come close to the nucleus are deflected by its positive charge. The few that head directly for the nucleus bounce back.
Niels Bohr refined Rutherford's model by proposing that electrons orbit the nucleus in fixed shells (energy levels) at specific distances. Electrons in higher shells have more energy. This model explained atomic emission spectra — when electrons fall from higher to lower energy levels they emit light at specific frequencies. Bohr's model is essentially the one used at GCSE.
The current scientific understanding replaces fixed orbits with probability clouds — regions where electrons are likely to be found. Neutrons were discovered by Chadwick in 1932, completing the picture of the nucleus. At GCSE, Bohr's shell model is sufficient.
The AQA atomic structure specification is at the AQA GCSE Chemistry specification page.
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