Atoms and Isotopes

Modern atomic model:

• Nucleus: central core containing positively charged protons and uncharged neutrons (collectively called nucleons).
• Electrons: negatively charged, orbit the nucleus in specific energy levels (shells).
• The atom is mostly empty space — the nucleus occupies a millionth of a billionth of the atom's volume.

Sizes:

• Atom diameter ≈ $10^{-10}$ m.
• Nucleus diameter ≈ $10^{-15}$ m (1 femtometre).
• If the atom were the size of a football stadium, the nucleus would be a pea on the centre spot.

Charge:

Particle	Charge	Relative mass
Proton	+1	1
Neutron	0	1
Electron	−1	~1/2000 (negligible)

Atom is neutral overall because number of protons = number of electrons.

Electrons exist in fixed energy levels (shells), not anywhere they like.

• An electron can ABSORB electromagnetic radiation (e.g. a photon of light) and jump to a HIGHER energy level.
• An electron in a high level can DROP back to a lower level by EMITTING electromagnetic radiation.

The frequency of the emitted radiation is set by the energy difference between the levels. This produces characteristic spectral lines (each element has its own 'fingerprint').

For GCSE, you don't need to calculate the frequencies — just understand that:

• Lower shells fill first (innermost shells have lowest energy).
• Atoms in the ground state have all electrons in the lowest available shells.
• An excited atom emits light when electrons return to lower shells.

Every nucleus is described by two numbers:

• Atomic number Z — number of protons. (Determines the element — Z = 1 means hydrogen, Z = 6 carbon, Z = 19 potassium.)
• Mass number A — total nucleons (protons + neutrons).

Standard notation:

$^A_Z X$

Example: $^{40}_{19}\text{K}$ means potassium with Z = 19 protons and A − Z = 40 − 19 = 21 neutrons.

Quick checks:

• Number of protons = Z.
• Number of neutrons = A − Z.
• Number of electrons in a neutral atom = Z.

Worked example. Carbon-14: $^{14}_6\text{C}$. 6 protons, 14 − 6 = 8 neutrons, 6 electrons (when neutral).

An isotope is an atom of an element with the same number of protons (so same Z) but a DIFFERENT number of neutrons (different A).

Examples:

• Hydrogen has three isotopes: $^1_1H$ (protium, 0 neutrons), $^2_1H$ (deuterium, 1 neutron), $^3_1H$ (tritium, 2 neutrons).
• Carbon: $^{12}_6C$ (most common, 6 neutrons), $^{13}_6C$ (7 neutrons, stable), $^{14}_6C$ (8 neutrons, radioactive — used for radiocarbon dating).
• Uranium: $^{235}_{92}U$ (143 n, used in fission), $^{238}_{92}U$ (146 n, most abundant).

Isotopes of the same element have identical chemistry (same electron arrangement) but different nuclear properties (mass, sometimes stability — radioactive isotopes covered in 4.4.2).

An ion is an atom (or molecule) with a net electrical charge — meaning the number of electrons no longer equals the number of protons.

• Positive ion (cation): lost one or more electrons (e.g. $\text{Na}^+$ has 11 protons, 10 electrons).
• Negative ion (anion): gained one or more electrons (e.g. $\text{Cl}^-$ has 17 protons, 18 electrons).

In radioactive decay (covered in 4.4.2), ions can form when alpha or beta particles ionise atoms in their path.

Some ions are formed by:

• Chemical reactions (e.g. NaCl in water).
• Light/UV/X-rays ejecting electrons (photoelectric effect).
• Friction/static electricity (rubbing).
• Radioactive particles colliding with atoms.

1. Dalton (1803). Atoms = tiny solid indivisible spheres. Different elements have different sphere types. (Atomic theory of matter.)

2. Thomson (1897). Discovered the electron (cathode ray experiments). Atoms can't be solid spheres — must contain negative charges. Proposed the plum-pudding model: a sphere of positive 'dough' with negative electrons embedded like raisins.

3. Rutherford (1909). Aimed alpha particles at a thin gold foil (Geiger and Marsden's experiment). EXPECTED most to pass through with small deflections (plum-pudding prediction). OBSERVED:

• Most went straight through.
• A few were deflected by large angles.
• A very few BOUNCED BACK.

Conclusion: the atom is mostly EMPTY SPACE (most pass through), with a small, dense, POSITIVELY CHARGED NUCLEUS (some α's bounce back — same-sign repulsion). This is the nuclear model.

4. Bohr (1913). Predicted electrons orbit only in specific energy levels. Solved the puzzle of why electrons don't spiral into the nucleus (classical physics predicted they should).

5. Chadwick (1932). Discovered the neutron — explained why the nucleus was heavier than the proton count alone predicted.

Geiger and Marsden (under Rutherford's direction) fired alpha particles (helium nuclei) at a thin gold foil. They detected where they ended up using a fluorescent screen.

Three observations:

1. Most α particles passed straight through with no deflection. → The atom is mostly EMPTY SPACE.

2. A small fraction were deflected through small angles. → They passed close to (but not directly into) something positive.

3. A very small number (1 in 8000) bounced back the way they came. → They hit a tiny, dense, positively-charged region — the nucleus.

This was incompatible with the plum-pudding model (which would have predicted only small deflections). The nuclear model replaced it.

Rutherford reportedly said the result was "almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you".

The atomic model has been revised repeatedly when experiments produced results the old model couldn't explain.

• Plum pudding couldn't explain alpha-scattering → nuclear model.
• Nuclear model couldn't explain electron stability → Bohr's energy levels.
• Bohr couldn't explain the mass difference (e.g. helium ≈ 4 amu but only 2 protons) → Chadwick added the neutron.

This pattern — evidence forces model change — is the scientific method in action. Modern physics still revises atomic models (quantum cloud model, Standard Model of particles), but the GCSE curriculum stops at Chadwick.