Atomic structure (C1.2)
Atomic structure underpins virtually every other chemistry topic. Expect short recall questions on subatomic particles and almost certainly a question on isotopes or electronic configuration.
Subatomic particles
| Particle | Symbol | Charge | Relative mass | Location |
|---|---|---|---|---|
| Proton | p | +1 | 1 | Nucleus |
| Neutron | n | 0 | 1 | Nucleus |
| Electron | e | −1 | negligible (1/1836) | Shells / orbitals around nucleus |
Atomic number and mass number
- Atomic number (Z) = number of protons in the nucleus. This defines the element — all atoms of an element have the same atomic number.
- Mass number A = total number of protons + neutrons in the nucleus.
- Number of neutrons = mass number − atomic number.
In a neutral atom, number of electrons = number of protons.
Example: sodium-23 (²³₁₁Na)
- Atomic number = 11 → 11 protons, 11 electrons.
- Mass number = 23 → neutrons = 23 − 11 = 12.
Electronic configuration
Electrons occupy energy levels (shells). The rules for filling are:
- 1st shell: max 2 electrons.
- 2nd shell: max 8 electrons.
- 3rd shell: max 8 electrons (at GCSE level).
Examples:
- Sodium (Z = 11): 2, 8, 1 (written as 2.8.1)
- Chlorine (Z = 17): 2, 8, 7
- Calcium (Z = 20): 2, 8, 8, 2
The number of electrons in the outer shell determines the chemical properties of the element and its group in the periodic table.
Group number = number of outer-shell electrons (for main-group elements).
Isotopes
Isotopes are atoms of the same element (same atomic number / same number of protons) with different numbers of neutrons (different mass numbers).
They have identical chemical properties (same electron configuration) but different physical properties (e.g. slightly different mass, different stability).
Examples:
- Carbon-12 (¹²C): 6 protons, 6 neutrons
- Carbon-13 (¹³C): 6 protons, 7 neutrons
- Carbon-14 (¹⁴C): 6 protons, 8 neutrons — radioactive; used in carbon dating
Chlorine exists as a mixture of ³⁵Cl (~75%) and ³⁷Cl (~25%). The relative atomic mass of 35.5 is a weighted average.
Relative atomic mass (Ar)
Ar = Σ (mass of isotope × abundance) / 100
Example — chlorine:
- Ar = (35 × 75 + 37 × 25) / 100 = (2625 + 925) / 100 = 3550 / 100 = 35.5
Development of the model of the atom
A brief history you may need to recall:
- Dalton (early 1800s) — atom as a solid sphere; smallest indivisible particle.
- Thomson (1897) — discovered the electron; proposed the plum-pudding model: electrons embedded in a positive sphere of charge.
- Rutherford (1911) — gold-foil scattering experiment: most particles passed straight through (atom is mostly empty space); a few deflected at large angles (positive nucleus is tiny and dense). Proposed the nuclear model.
- Bohr (1913) — electrons orbit in defined shells (energy levels) at fixed distances from the nucleus.
- Modern model — electrons are in probability orbitals around a nucleus of protons and neutrons.
⚠ OCR loves asking: "What did the Rutherford scattering experiment show that the plum-pudding model could not explain?" Answer: it showed that most mass and all positive charge is concentrated in a tiny nucleus, not spread out; the plum-pudding model predicted no large deflections.
Common Gateway-paper mistakes
- Confusing atomic number (protons) with mass number (protons + neutrons).
- Forgetting that electrons have negligible mass — they don't contribute to the mass number.
- Saying isotopes have different chemical properties — they have the SAME chemical properties.
- Miscounting electrons from the outer shell when writing electronic configurations.
- Mixing up Ar with mass number (Ar is a weighted average; mass number is always a whole number for a specific isotope).
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