What you'll learn
This topic forms the foundation of atomic structure in AQA GCSE Chemistry. You'll learn about the three subatomic particles that make up atoms, understand how to interpret atomic number and mass number from the periodic table, and calculate the numbers of protons, neutrons and electrons in any atom or ion. These concepts appear throughout Paper 1 and Paper 2, particularly in questions about the periodic table, bonding, and chemical reactions.
Key terms and definitions
Atom — the smallest particle of an element that can exist, consisting of a nucleus containing protons and neutrons, surrounded by electrons in shells.
Proton — a positively charged subatomic particle found in the nucleus, with a relative mass of 1 and relative charge of +1.
Neutron — a neutral subatomic particle found in the nucleus, with a relative mass of 1 and relative charge of 0.
Electron — a negatively charged subatomic particle found in shells (energy levels) around the nucleus, with a relative mass of approximately 1/2000 (often taken as 0) and relative charge of −1.
Atomic number — the number of protons in the nucleus of an atom; this defines the element and is represented by the symbol Z.
Mass number — the total number of protons and neutrons in the nucleus of an atom, represented by the symbol A.
Ion — a charged particle formed when an atom loses or gains electrons; positive ions (cations) have fewer electrons than protons, while negative ions (anions) have more electrons than protons.
Isotope — atoms of the same element (same atomic number) with different numbers of neutrons, and therefore different mass numbers.
Core concepts
The structure of the atom
Atoms consist of a central nucleus surrounded by electrons. The nucleus is extremely small compared to the overall size of the atom — if an atom were the size of a football stadium, the nucleus would be smaller than a pea at the centre.
The nucleus contains:
- Protons (positive charge)
- Neutrons (no charge)
Around the nucleus:
- Electrons orbit in shells or energy levels (negative charge)
The nucleus contains most of the atom's mass because protons and neutrons each have a relative mass of 1, while electrons have negligible mass. Despite being tiny, the nucleus is extremely dense.
In a neutral atom, the number of protons equals the number of electrons. This means the positive charges balance the negative charges, giving an overall charge of zero.
Relative charges and masses of subatomic particles
You must know these values for AQA GCSE Chemistry exams:
| Particle | Relative mass | Relative charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/2000 (or 0) | −1 | Shells around nucleus |
The relative mass of an electron is so small (approximately 0.0005) that it's often recorded as zero or 1/2000 in GCSE calculations. This means virtually all of an atom's mass comes from the protons and neutrons in its nucleus.
The charges balance in a neutral atom: if there are 6 protons (+6), there must be 6 electrons (−6), giving an overall charge of zero.
Atomic number and mass number
Every element has a unique atomic number (Z), which tells you the number of protons in the nucleus. This number defines what element an atom is — you cannot change the number of protons without changing the element itself.
The mass number (A) tells you the total number of protons plus neutrons in the nucleus.
These numbers are shown in standard notation:
- Mass number (A) is written at the top left of the element symbol
- Atomic number (Z) is written at the bottom left of the element symbol
Example: ²³₁₁Na
- Mass number = 23
- Atomic number = 11
From this information you can calculate:
- Number of protons = atomic number = 11
- Number of electrons (in a neutral atom) = atomic number = 11
- Number of neutrons = mass number − atomic number = 23 − 11 = 12
Using the periodic table
On the periodic table, each element box contains:
- The element symbol (e.g., C for carbon)
- The atomic number (smaller number, usually at the top)
- The relative atomic mass (larger number, usually at the bottom)
The atomic number is always a whole number and tells you the number of protons. For GCSE calculations involving specific atoms, you'll be given the mass number directly in the question, as the relative atomic mass on the periodic table is an average that accounts for different isotopes.
Remember: the number in the name of an isotope is its mass number. For example, carbon-12 has a mass number of 12, and carbon-14 has a mass number of 14. Both have 6 protons (because they're both carbon), but carbon-12 has 6 neutrons while carbon-14 has 8 neutrons.
Calculating numbers of subatomic particles
For any atom, you can calculate the three subatomic particle numbers using these rules:
For neutral atoms:
- Number of protons = atomic number
- Number of electrons = atomic number (because charges balance)
- Number of neutrons = mass number − atomic number
For ions:
- Number of protons = atomic number (this never changes)
- Number of neutrons = mass number − atomic number (this never changes unless it's a different isotope)
- Number of electrons = atomic number − charge on the ion
Examples of ion calculations:
- Na⁺ (sodium ion with +1 charge): sodium has atomic number 11, so the Na⁺ ion has 11 protons but only 10 electrons (it has lost 1 electron)
- O²⁻ (oxide ion with −2 charge): oxygen has atomic number 8, so the O²⁻ ion has 8 protons but 10 electrons (it has gained 2 electrons)
- Al³⁺ (aluminium ion with +3 charge): aluminium has atomic number 13, so the Al³⁺ ion has 13 protons but only 10 electrons (it has lost 3 electrons)
The key rule: positive ions have lost electrons (fewer electrons than protons), while negative ions have gained electrons (more electrons than protons).
Isotopes
Isotopes are different forms of the same element. They have:
- The same atomic number (same number of protons)
- Different mass numbers (different numbers of neutrons)
- Identical chemical properties (because chemical reactions involve electrons, and isotopes have the same electron arrangement)
- Slightly different physical properties (because they have different masses)
Common examples you should know:
Chlorine isotopes:
- Chlorine-35: ³⁵₁₇Cl (17 protons, 18 neutrons, 17 electrons)
- Chlorine-37: ³⁷₁₇Cl (17 protons, 20 neutrons, 17 electrons)
Carbon isotopes:
- Carbon-12: ¹²₆C (6 protons, 6 neutrons, 6 electrons)
- Carbon-13: ¹³₆C (6 protons, 7 neutrons, 6 electrons)
- Carbon-14: ¹⁴₆C (6 protons, 8 neutrons, 6 electrons) — radioactive isotope used in carbon dating
Most elements exist as a mixture of isotopes in nature. The relative atomic mass shown on the periodic table is a weighted average of all the naturally occurring isotopes. This is why relative atomic masses are rarely whole numbers (e.g., chlorine's relative atomic mass is 35.5 because it's roughly 75% chlorine-35 and 25% chlorine-37).
Worked examples
Example 1: Calculating subatomic particles in a neutral atom
Question: An atom of fluorine is represented as ¹⁹₉F. State the number of protons, neutrons and electrons in this atom. [3 marks]
Solution:
- Number of protons = atomic number = 9 ✓
- Number of neutrons = mass number − atomic number = 19 − 9 = 10 ✓
- Number of electrons = atomic number = 9 (in a neutral atom) ✓
Mark scheme note: One mark for each correct answer. Make sure you show your working for the neutron calculation.
Example 2: Comparing isotopes
Question: Magnesium has three isotopes: magnesium-24, magnesium-25 and magnesium-26. The atomic number of magnesium is 12.
(a) Explain what is meant by the term isotope. [2 marks] (b) Complete the table below for these three isotopes. [3 marks]
| Isotope | Number of protons | Number of neutrons | Number of electrons |
|---|---|---|---|
| Mg-24 | |||
| Mg-25 | |||
| Mg-26 |
Solution:
(a) Isotopes are atoms of the same element ✓ that have the same number of protons but different numbers of neutrons (or different mass numbers). ✓
Mark scheme note: Must mention "same element" or "same number of protons" for first mark, and "different numbers of neutrons" or "different mass numbers" for second mark.
(b)
| Isotope | Number of protons | Number of neutrons | Number of electrons |
|---|---|---|---|
| Mg-24 | 12 | 12 | 12 |
| Mg-25 | 12 | 13 | 12 |
| Mg-26 | 12 | 14 | 12 |
✓ All protons correct (all 12) ✓ All neutrons correct (12, 13, 14) ✓ All electrons correct (all 12)
Mark scheme note: Three marks awarded for completely correct table. Partial marks may be given if only one row or column contains errors.
Example 3: Ions and subatomic particles
Question: A calcium ion has the symbol Ca²⁺. The atomic number of calcium is 20 and the mass number of this particular calcium atom is 40.
State the number of protons, neutrons and electrons in this calcium ion. [3 marks]
Solution:
- Number of protons = atomic number = 20 ✓
- Number of neutrons = mass number − atomic number = 40 − 20 = 20 ✓
- Number of electrons = 20 − 2 = 18 (the ion has a 2+ charge, meaning it has lost 2 electrons) ✓
Mark scheme note: Common error is to give 20 electrons. The 2+ charge indicates the ion has lost 2 electrons, so must have 18 electrons remaining.
Common mistakes and how to avoid them
Confusing mass number and atomic number — Remember: atomic number is always smaller and equals the number of protons. Mass number is larger and equals protons plus neutrons. On the periodic table, the atomic number is the defining characteristic of each element.
Forgetting that ions have different numbers of electrons — In ions, the number of protons never changes, but the number of electrons does. Positive ions have lost electrons (fewer electrons than protons), negative ions have gained electrons (more electrons than protons). The number of neutrons also stays the same unless you're dealing with different isotopes.
Mixing up the positions of mass and atomic numbers in notation — In standard notation like ²³₁₁Na, the mass number (23) is always at the top, and the atomic number (11) is always at the bottom. A useful memory aid: mass is massive, so it goes on top.
Thinking isotopes have different numbers of electrons — Isotopes differ only in their number of neutrons. They have the same number of protons and (if they're neutral atoms) the same number of electrons. This is why isotopes have identical chemical properties.
Assuming all atoms of an element have the same mass number — Different isotopes exist for most elements. When you're given a specific mass number in a question (e.g., carbon-14), use that specific value. Don't use the relative atomic mass from the periodic table for subatomic particle calculations.
Not showing working for neutron calculations — Always write the calculation: neutrons = mass number − atomic number. Even if you can do it mentally, showing your working can earn method marks if you make an arithmetic error.
Exam technique for "Subatomic particles, atomic number and mass number"
Command word "State" — Give a brief answer without explanation. For example, "State the number of protons" requires just the number, not a calculation or explanation. However, showing your working for neutrons is still recommended as it can earn method marks.
Command word "Explain" — You must give reasons or make relationships clear. For isotope questions, stating that isotopes have "different mass numbers" alone won't earn full marks — you need to explain this is because they have different numbers of neutrons.
Reading notation carefully — If you see ³⁵₁₇Cl or Cl-35, both refer to chlorine-35. Extract the atomic number (17) and mass number (35) before attempting calculations. Check whether you're dealing with an atom or an ion (look for + or − symbols).
Marks allocation guides your answer — A 3-mark question about subatomic particles typically wants three numbers (protons, neutrons, electrons). A 2-mark definition of isotopes requires two distinct points. Use the marks available to check you've given sufficient detail.
Quick revision summary
Atoms contain three subatomic particles: protons (+1 charge, mass 1) and neutrons (0 charge, mass 1) in the nucleus, plus electrons (−1 charge, negligible mass) in shells around the nucleus. The atomic number equals the number of protons and defines the element. The mass number equals protons plus neutrons. In neutral atoms, electrons equal protons. Ions form when atoms lose or gain electrons. Isotopes are atoms of the same element with different numbers of neutrons, giving different mass numbers but identical chemical properties.