CC1.2 — Atomic structure (Edexcel 1SC0)
Subatomic particles
| Particle | Location | Relative mass | Relative charge |
|---|---|---|---|
| Proton | Nucleus | 1 | +1 |
| Neutron | Nucleus | 1 | 0 |
| Electron | Shells | 1/1836 (~0) | −1 |
- Atomic number (Z): number of protons. This defines the element.
- Mass number A: protons + neutrons.
- Number of neutrons = A − Z.
- In a neutral atom: number of protons = number of electrons.
Isotopes
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
Example: Carbon-12 (6p, 6n) and Carbon-14 (6p, 8n). Same chemical properties (same electron configuration); different mass.
Relative atomic mass (Ar): weighted mean mass of all isotopes, relative to carbon-12.
Development of the atomic model
- Dalton (early 1800s): atoms are solid spheres, different for each element.
- Thomson (1897): discovered electron; proposed the "plum pudding" model — electrons embedded in a positive sphere.
- Rutherford (1909): gold foil experiment — most alpha particles passed through; some deflected; nuclear model proposed — small, dense, positive nucleus surrounded by electrons.
- Bohr (1913): electrons in discrete energy levels (shells).
- Chadwick (1932): discovered the neutron.
AI-generated · claude-opus-4-7 · v3-edexcel-combined-science