At GCSE, the atom goes beyond protons, neutrons, and electrons in shells: you need to explain isotopes, calculate relative atomic mass from isotope data, write electron configurations up to calcium, and explain how atoms form ions. All of this flows from one model — the nuclear atom, where almost all the mass sits in a tiny positive nucleus.
What does GCSE atomic structure add to the KS3 model?
At KS3, atomic structure is taught as: a positive nucleus containing protons and neutrons, surrounded by electrons in shells. At GCSE, the detail deepens:
- You must know the relative masses and charges of subatomic particles
- You must use atomic number and mass number to calculate numbers of each particle
- You must understand isotopes and calculate relative atomic mass from percentage abundance data
- You must write electron configurations for elements up to calcium (Z = 20)
- You must explain ion formation in terms of electrons gained or lost
| Particle | Relative mass | Relative charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/1836 (effectively 0) | −1 | Shells around nucleus |
How do you use atomic number and mass number?
The atomic number (Z) = number of protons. In a neutral atom, this equals the number of electrons.
The mass number (A) = number of protons + number of neutrons.
Therefore: number of neutrons = mass number − atomic number
Example: Sodium is written as ²³₁₁Na
- Atomic number = 11 → 11 protons, 11 electrons (neutral atom)
- Mass number = 23 → 23 − 11 = 12 neutrons
| Element | Symbol | Atomic number (Z) | Mass number (A) | Protons | Neutrons | Electrons |
|---|---|---|---|---|---|---|
| Hydrogen | ¹₁H | 1 | 1 | 1 | 0 | 1 |
| Carbon | ¹²₆C | 6 | 12 | 6 | 6 | 6 |
| Oxygen | ¹⁶₈O | 8 | 16 | 8 | 8 | 8 |
| Chlorine | ³⁵₁₇Cl | 17 | 35 | 17 | 18 | 17 |
| Calcium | ⁴⁰₂₀Ca | 20 | 40 | 20 | 20 | 20 |
What are isotopes?
Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Because the number of protons defines the element, isotopes have identical chemical properties (same electron configuration → same reactions). They differ only in mass.
Examples:
| Isotope | Protons | Neutrons | Symbol |
|---|---|---|---|
| Carbon-12 | 6 | 6 | ¹²₆C |
| Carbon-13 | 6 | 7 | ¹³₆C |
| Carbon-14 | 6 | 8 | ¹⁴₆C |
| Chlorine-35 | 17 | 18 | ³⁵₁₇Cl |
| Chlorine-37 | 17 | 20 | ³⁷₁₇Cl |
Carbon-12 and carbon-14 participate in the same chemical reactions (they react identically with oxygen to form CO₂), but carbon-14 is radioactive because of its unstable nucleus — a property that depends on the nucleus, not the electrons.
How do you calculate relative atomic mass from isotope data?
Relative atomic mass (Ar) is the weighted average mass of all isotopes of an element, taking into account their percentage abundance in nature.
$$A_r = \frac{(\text{mass of isotope 1} \times % \text{ abundance}_1) + (\text{mass of isotope 2} \times % \text{ abundance}_2)}{100}$$
Worked example — chlorine:
Chlorine exists as two isotopes: 75% Cl-35 and 25% Cl-37.
- A_r = (35 × 75 + 37 × 25) / 100
- A_r = (2,625 + 925) / 100
- A_r = 3,550 / 100 = 35.5
This explains why the periodic table shows 35.5 for chlorine rather than a whole number — it is an average weighted by natural abundance.
How do you write electron configurations?
Electrons occupy shells (energy levels) around the nucleus. The rules for the first three shells are:
- Shell 1: maximum 2 electrons
- Shell 2: maximum 8 electrons
- Shell 3: maximum 8 electrons (at GCSE level, up to calcium)
Electrons fill the lowest available shell first. Write configurations as a sequence of numbers (one per shell):
| Element | Z | Electron configuration |
|---|---|---|
| Hydrogen | 1 | 1 |
| Helium | 2 | 2 |
| Lithium | 3 | 2, 1 |
| Carbon | 6 | 2, 4 |
| Nitrogen | 7 | 2, 5 |
| Oxygen | 8 | 2, 6 |
| Fluorine | 9 | 2, 7 |
| Neon | 10 | 2, 8 |
| Sodium | 11 | 2, 8, 1 |
| Magnesium | 12 | 2, 8, 2 |
| Phosphorus | 15 | 2, 8, 5 |
| Sulfur | 16 | 2, 8, 6 |
| Argon | 18 | 2, 8, 8 |
| Potassium | 19 | 2, 8, 8, 1 |
| Calcium | 20 | 2, 8, 8, 2 |
The electron configuration determines chemical properties: elements with the same number of electrons in their outer shell (same group in the periodic table) have similar reactions.
How do atoms form ions?
Ions are formed when atoms gain or lose electrons, giving them an overall electrical charge.
- Losing electrons → positive ion (cation). Metals in Groups 1, 2, and 3 typically lose electrons.
- Gaining electrons → negative ion (anion). Non-metals in Groups 6 and 7 typically gain electrons.
The number of electrons gained or lost equals the group number distance from the nearest noble gas configuration.
Examples:
| Atom | Change | Ion formed | Why |
|---|---|---|---|
| Na (2,8,1) | Loses 1 e⁻ | Na⁺ (2,8) | Achieves stable neon configuration |
| Mg (2,8,2) | Loses 2 e⁻ | Mg²⁺ (2,8) | Achieves stable neon configuration |
| Cl (2,8,7) | Gains 1 e⁻ | Cl⁻ (2,8,8) | Achieves stable argon configuration |
| O (2,6) | Gains 2 e⁻ | O²⁻ (2,8) | Achieves stable neon configuration |
The proton number is unchanged when ions form — sodium is always sodium (11 protons). Only the electron number changes.
Frequently asked questions
Why do isotopes of the same element have the same chemical properties?
Chemical properties are determined by the electron configuration — specifically the number of electrons in the outer shell. Isotopes differ only in their number of neutrons, which affects nuclear stability and mass but not electron configuration. Because all isotopes of chlorine have 17 protons and therefore 17 electrons in the same configuration (2,8,7), they all react identically — gaining one electron to form Cl⁻.
What is the difference between relative atomic mass and mass number?
The mass number is a whole number giving the total number of protons and neutrons in one specific isotope (e.g. C-12 has mass number 12 exactly). The relative atomic mass is the weighted average of all naturally occurring isotopes and is therefore often not a whole number (e.g. chlorine is 35.5; bromine is 79.9). The relative atomic mass is the value shown on the periodic table.
How do I calculate the number of neutrons when the element has a non-integer relative atomic mass?
Use the mass number of a specific isotope, not the relative atomic mass. For example, for Cl-35: neutrons = 35 − 17 = 18. For Cl-37: neutrons = 37 − 17 = 20. The relative atomic mass of 35.5 is an average and does not represent any single atom; no chlorine atom has 18.5 neutrons.
Why does the arrangement of electrons in shells explain the periodic table's structure?
The period number equals the number of occupied electron shells. The group number equals the number of electrons in the outer shell (for Groups 1–7) or tells you the element has a full outer shell (Group 0/8). Elements in the same group have the same outer-shell electron count, explaining why they show similar chemical reactions — lithium, sodium, and potassium all have one outer electron and all react vigorously with water, for instance.
For Socratic GCSE chemistry with Professor Curie — building from sub-atomic particles to the patterns of the periodic table — visit aitutors.me.