The Atom and Subatomic Particles

The atom

Everything is made of atoms. An atom has a tiny central nucleus containing protons and neutrons, surrounded by electrons arranged in shells (energy levels). Atoms are incredibly small — the radius is about 0.1 nanometres (1 × 10⁻¹⁰ m) and the nucleus is around 1/10 000 of that.

The three subatomic particles

Particle	Relative charge	Relative mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	−1	very small (~1/1836)	Shells around the nucleus

• Atoms have no overall charge because the number of protons = number of electrons (+1 and −1 cancel).
• Almost all the mass is in the nucleus (protons + neutrons); electrons have negligible mass.

Atomic number and mass number

• Atomic number = number of protons (this defines the element and its place in the periodic table).
• Mass number = number of protons + neutrons.
• Number of neutrons = mass number − atomic number.

For example, sodium (Na) has atomic number 11 and mass number 23 → 11 protons, 11 electrons, and 23 − 11 = 12 neutrons.

Electron shells (energy levels)

Electrons fill shells from the inside out:

• 1st shell holds up to 2 electrons.
• 2nd and 3rd shells hold up to 8 each.

Example — sodium (11 electrons): 2, 8, 1. The number of electrons in the outer shell matches the group in the periodic table.

Isotopes

Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons — so the same atomic number but a different mass number. Example: chlorine-35 and chlorine-37.

• Relative atomic mass (Aᵣ) is the average mass of an element's atoms, taking isotope abundance into account (which is why chlorine's Aᵣ is 35.5, not a whole number).

History of the atom (models)

Ideas changed as evidence grew: Dalton (solid spheres) → Thomson (plum pudding) → Rutherford (nuclear model, from the alpha-scattering experiment) → Bohr (electrons in shells). This shows how scientific models are revised with new evidence.

Worked example

An atom has atomic number 17 and mass number 35. Give its particle counts.

• Protons = 17, Electrons = 17, Neutrons = 35 − 17 = 18. ✓

Common mistakes

• Forgetting mass number − atomic number = neutrons.
• Saying isotopes have different protons — they differ only in neutrons.
• Filling electron shells incorrectly (remember 2, then 8, then 8).

Exam tips

• Learn the charge and mass of each particle — a guaranteed easy mark.
• Practise working out protons/neutrons/electrons from atomic and mass numbers.
• Explain isotopes clearly and link them to non-whole-number relative atomic masses.

Key facts to remember

• Atom = nucleus (protons + neutrons) + electrons in shells; overall neutral (protons = electrons).
• Atomic number = protons; mass number = protons + neutrons; neutrons = mass − atomic number.
• Isotopes = same protons, different neutrons; electron shells fill 2, 8, 8.