GCSE Physics revision

Mass number, atomic number and isotopes

Atoms and isotopes

AQA 4.4.1.2

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AQA student objectives

Learning pathway

ALL 🎯

Relate differences between isotopes to differences in conventional representations of their identities, charges and masses.

MOST 🎯🎯

Apply the specified scientific knowledge of mass number, atomic number and isotopes to a relevant example.

SOME 🎯🎯🎯

Analyse an unfamiliar example of mass number, atomic number and isotopes using the specified scientific ideas.

Revision summary

Read on screen, then print for Cornell-style active revision.

Atoms are incredibly tiny, with a radius of approximately 1 imes 10^{-10} metres (1 tenbillionth of a metre).
To appreciate the scale, there are more atoms in a single grain of sand than there are grains of sand on every beach on Earth.

The nucleus sits at the centre of the atom, is positively charged, and contains the heavy particles — protons and neutrons.
Electrons are tiny, negatively charged particles that orbit the nucleus in shells at various energy levels.
An atom is mostly empty space — if an atom were the size of a football stadium, the nucleus would be roughly the size of a pea on the centre spot.

2. Structure of the Modern AtomAQA

The nucleus sits at the centre of the atom, is positively charged, and contains the heavy particles — protons and neutrons.
Electrons are tiny, negatively charged particles that orbit the nucleus in shells at various energy levels.
An atom is mostly empty space — if an atom were the size of a football stadium, the nucleus would be roughly the size of a pea on the centre spot.

Protons carry a relative charge of +1 and a relative mass of 1.
Neutrons carry no charge (0) and have a relative mass of 1.
Electrons carry a relative charge of −1 but have a negligible mass compared to protons and neutrons.

The atomic number is the number of protons in the nucleus and uniquely identifies the element — for example, every carbon atom has exactly 6 protons.
The mass number is the total number of protons and neutrons in the nucleus, calculated as: Mass number = protons + neutrons.
To find the number of neutrons, subtract the atomic number from the mass number: neutrons = mass number - atomic number.

4. Atomic Number and Mass NumberAQA

The atomic number is the number of protons in the nucleus and uniquely identifies the element — for example, every carbon atom has exactly 6 protons.
The mass number is the total number of protons and neutrons in the nucleus, calculated as: Mass number = protons + neutrons.
To find the number of neutrons, subtract the atomic number from the mass number: neutrons = mass number - atomic number.

Atoms are electrically neutral, meaning the number of negatively charged electrons always equals the number of positively charged protons.
This perfect balance of charges means the overall charge of a neutral atom is zero.

Isotopes are atoms of the same element that have the same number of protons (same atomic number) but a different number of neutrons.
Because isotopes differ only in neutron number, they have different mass numbers but identical chemical properties.
Carbon-12 and Carbon-14 are classic examples of isotopes — both have 6 protons, but Carbon-14 has 2 extra neutrons.

6. IsotopesAQA

Isotopes are atoms of the same element that have the same number of protons (same atomic number) but a different number of neutrons.
Because isotopes differ only in neutron number, they have different mass numbers but identical chemical properties.
Carbon-12 and Carbon-14 are classic examples of isotopes — both have 6 protons, but Carbon-14 has 2 extra neutrons.

In the early 1800s, atoms were thought to be tiny, solid, indivisible spheres — the 'billiard ball' model.
The discovery of the electron led to the plum pudding model, where negative electrons were embedded in a ball of diffuse positive charge.
In 1911, Rutherford's nuclear model replaced the plum pudding model, proposing a tiny, dense, positively charged nucleus at the centre.
Bohr refined the nuclear model shortly after by suggesting electrons orbit the nucleus in fixed shells at specific energy levels.
In 1932, Chadwick discovered the neutron, completing our understanding of the nucleus and explaining why the mass number can differ from the atomic number.

7. The History of Atomic ModelsAQA

In the early 1800s, atoms were thought to be tiny, solid, indivisible spheres — the 'billiard ball' model.
The discovery of the electron led to the plum pudding model, where negative electrons were embedded in a ball of diffuse positive charge.
In 1911, Rutherford's nuclear model replaced the plum pudding model, proposing a tiny, dense, positively charged nucleus at the centre.
Bohr refined the nuclear model shortly after by suggesting electrons orbit the nucleus in fixed shells at specific energy levels.
In 1932, Chadwick discovered the neutron, completing our understanding of the nucleus and explaining why the mass number can differ from the atomic number.

In Rutherford's famous experiment, alpha particles were fired at an ultra-thin sheet of gold foil to probe the structure of the atom.
Most alpha particles passed straight through the foil, suggesting that atoms are mostly empty space.
A small number of alpha particles were deflected at large angles, and some bounced almost straight back — a result that was completely unexpected.
The only explanation for this was that the atom's positive charge and most of its mass must be concentrated in a tiny, dense nucleus — disproving the plum pudding model.

8. The Alpha Particle Scattering ExperimentAQA

In Rutherford's famous experiment, alpha particles were fired at an ultra-thin sheet of gold foil to probe the structure of the atom.
Most alpha particles passed straight through the foil, suggesting that atoms are mostly empty space.
A small number of alpha particles were deflected at large angles, and some bounced almost straight back — a result that was completely unexpected.
The only explanation for this was that the atom's positive charge and most of its mass must be concentrated in a tiny, dense nucleus — disproving the plum pudding model.