Atomic Structure

This lesson covers the structure of the atom and the sub-atomic particles within it, as required by AQA GCSE Chemistry specification (5.1.1). Understanding atomic structure is fundamental to chemistry — it explains why elements behave the way they do, how they bond, and why the periodic table is arranged the way it is.

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The Structure of an Atom

An atom consists of a central nucleus surrounded by electrons that orbit in shells (energy levels). The nucleus is extremely small compared to the overall size of the atom.

Part of the Atom	Description
Nucleus	Found at the centre of the atom. Contains protons and neutrons. Very small and dense. Positively charged overall.
Electron shells	Regions around the nucleus where electrons are found. Electrons orbit the nucleus at specific energy levels (shells).

The radius of an atom is approximately 1 x 10^-10 m (0.1 nm). The radius of the nucleus is about 1 x 10^-14 m — roughly 10,000 times smaller than the atom itself. This means that most of an atom is empty space.

Exam Tip: The nucleus is approximately 10,000 times smaller than the atom. If the atom were the size of a football stadium, the nucleus would be the size of a marble on the centre spot. This analogy helps you appreciate how much of an atom is empty space.

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Sub-Atomic Particles

Atoms are made up of three types of sub-atomic particles:

Particle	Relative Charge	Relative Mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	Very small (1/1836, negligible)	Shells around the nucleus

Key Points

• Protons and neutrons have approximately the same mass — we call this a relative mass of 1.
• Electrons have a negligibly small mass (about 1/1836 of a proton).
• The mass of an atom is concentrated almost entirely in the nucleus.
• In a neutral atom, the number of protons equals the number of electrons, so the charges balance out and the atom has no overall charge.

Exam Tip: Electrons have virtually zero mass compared to protons and neutrons. When calculating the mass of an atom, you can ignore the mass of electrons. However, you must NEVER ignore their charge — electrons are essential for determining the charge of ions.

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Atomic Number and Mass Number

Every element is defined by the number of protons in its nucleus. Two key numbers describe each atom:

Term	Symbol	Definition
Atomic number (proton number)	Z	The number of protons in the nucleus. This defines which element the atom is.
Mass number (nucleon number)	A	The total number of protons + neutrons in the nucleus.

Calculating Sub-Atomic Particles

• Number of protons = atomic number
• Number of electrons = atomic number (in a neutral atom)
• Number of neutrons = mass number - atomic number

Example: Sodium (Na)

• Atomic number = 11, Mass number = 23
• Protons = 11
• Electrons = 11
• Neutrons = 23 - 11 = 12

Example: Carbon (C)

• Atomic number = 6, Mass number = 12
• Protons = 6
• Electrons = 6
• Neutrons = 12 - 6 = 6

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Atomic Structure Diagram

graph TD
    A["Atom"] --> B["Nucleus"]
    A --> C["Electron Shells"]
    B --> D["Protons (+1 charge, mass 1)"]
    B --> E["Neutrons (0 charge, mass 1)"]
    C --> F["Electrons (-1 charge, negligible mass)"]
    A --> G["Atomic Number (Z) = number of protons"]
    A --> H["Mass Number (A) = protons + neutrons"]
    A --> I["Neutral atom: protons = electrons"]

    style A fill:#2c3e50,color:#fff
    style B fill:#e74c3c,color:#fff
    style C fill:#3498db,color:#fff
    style D fill:#e74c3c,color:#fff
    style E fill:#95a5a6,color:#fff
    style F fill:#3498db,color:#fff

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Isotopes

Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons. This means they have the same atomic number but different mass numbers.

Isotope	Protons	Neutrons	Mass Number
Carbon-12	6	6	12
Carbon-13	6	7	13
Carbon-14	6	8	14
Chlorine-35	17	18	35
Chlorine-37	17	20	37

Properties of Isotopes

• Isotopes of the same element have identical chemical properties because they have the same number of electrons (and it is electrons that determine chemical behaviour).
• Isotopes have different physical properties (such as density and rate of diffusion) because they have different masses.
• Some isotopes are radioactive (unstable nuclei that decay by emitting radiation) — these are called radioisotopes.

Exam Tip: The question "Explain why isotopes of an element have the same chemical properties" is very common. The answer is always: they have the same number of electrons in the same arrangement, and it is the electron configuration that determines chemical properties.

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Relative Atomic Mass (Ar)

Because most elements exist as a mixture of isotopes, the relative atomic mass (Ar) of an element is a weighted average of the masses of all its isotopes, taking into account their relative abundances.

The formula is:

Ar = sum of (isotope mass x percentage abundance) / 100

Example: Chlorine

Chlorine has two isotopes:

• Chlorine-35 (75% abundance)
• Chlorine-37 (25% abundance)

Ar = (35 x 75 + 37 x 25) / 100 Ar = (2625 + 925) / 100 Ar = 3550 / 100 Ar = 35.5

This is why the relative atomic mass of chlorine on the periodic table is 35.5, not a whole number.

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The Development of the Atomic Model

The model of the atom has changed over time as scientists have made new discoveries. You need to know the key stages:

Scientist / Model	Date	Key Idea
John Dalton	Early 1800s	Atoms are tiny, indivisible solid spheres. Each element has its own type of atom.
J.J. Thomson	1897	Discovered the electron. Proposed the "plum pudding" model — a ball of positive charge with electrons embedded in it.
Ernest Rutherford	1909	Alpha particle scattering experiment showed that most of the atom is empty space with a small, dense, positively charged nucleus. Proposed the nuclear model.
Niels Bohr	1913	Electrons orbit the nucleus in specific shells (energy levels), not randomly. Each shell has a fixed energy.
James Chadwick	1932	Discovered the neutron — a neutral particle in the nucleus with the same mass as a proton.

Rutherford's Alpha Particle Scattering Experiment

Rutherford, Geiger and Marsden fired alpha particles (positively charged) at a thin sheet of gold foil and observed: