Atomic Structure, Protons, Neutrons and Isotopes
Everything in Chemistry starts here. One confusion dominates the whole topic — mixing up the two numbers on the periodic table — and it was recorded more often than any other error in the corpus behind these notes.
1. The subatomic particles
| Particle | Location | Relative charge | Relative mass |
|---|---|---|---|
| Proton | nucleus | +1 | 1 |
| Neutron | nucleus | 0 | 1 |
| Electron | shells around the nucleus | −1 | 1/1840 (negligible) |
A neutron is neutral; an electron is negative. Confusing their charges was recorded — the clue is in the name neutron.
Almost all the mass is in the NUCLEUS, because protons and neutrons have mass 1 each and electrons have virtually none.
But protons alone do not account for the mass — neutrons count too. A recorded error assumed only protons contribute. The mass number is protons plus neutrons.
Electrons are much smaller and lighter than protons and neutrons, which are roughly equal to each other.
Atoms cannot be seen with an ordinary microscope. A recorded error claimed some can — atoms are far too small for visible light.
2. The two numbers — get these right
Proton number (atomic number, Z) = number of protons. It is the SMALLER number, and it identifies the element. Nucleon number (mass number, A) = protons + neutrons. It is the LARGER number.
Number of NEUTRONS = nucleon number − proton number = A − Z.
This is the single most important calculation in the topic, and confusing the two numbers was recorded more than a dozen times.
A neutral atom has equal protons and electrons, so the charges cancel.
Worked example — ⁶⁵₃₀Zn:
- Protons = 30
- Electrons = 30 (neutral atom)
- Neutrons = 65 − 30 = 35
Check the smaller number is the proton number. If your “neutron” count comes out negative, you have subtracted the wrong way round.
Use the periodic table to check. Tutors advised this directly — look up the element and confirm both numbers. Several recorded errors were simple misidentifications: sulfur for oxygen, silicon for sulfur, sulfur for hydrogen.
Nucleon number and mass number are the same thing — two names for A.
3. Isotopes
Isotopes are atoms of the SAME element with the SAME number of protons but DIFFERENT numbers of neutrons.
This definition is worth 2 marks — tutors flagged it explicitly. You must state both halves: same protons and different neutrons.
Example: chlorine-35 (17p, 18n) and chlorine-37 (17p, 20n). Both have 17 protons, so both are chlorine.
Isotopes must have the SAME proton number. A recorded error had isotopes with different proton numbers — but changing the protons changes the element entirely.
Isotopes differ in NEUTRONS, not electrons. Changing the electrons makes an ion, not an isotope.
Properties:
Isotopes have identical CHEMICAL properties, because chemical behaviour depends on the electronic configuration — and isotopes of an element have the same number of electrons, arranged identically. They have different PHYSICAL properties (mass, density, rate of diffusion), because they have different masses.
Say “same electronic configuration” when explaining why chemical properties match. Confusion between isotopes and electronic configuration was recorded repeatedly — the point is that they are the same, which is exactly why the chemistry is identical.
4. Ions
An ion is an atom (or group of atoms) that has lost or gained electrons, giving it a charge.
- Losing electrons → positive ion (cation) — metals
- Gaining electrons → negative ion (anion) — non-metals
The number of PROTONS never changes when an ion forms — only the electrons. So the proton number and nucleon number are unchanged.
Worked example — Al³⁺:
- Aluminium has proton number 13
- The 3+ charge means it has lost 3 electrons
- So it has 13 protons and 10 electrons
Electronic configuration of Al³⁺ is 2,8 — not 2,8,3. This exact error was recorded twice. Work out the configuration after removing the electrons.
Writing charges:
Write the number BEFORE the sign for an ion charge: 2+, 3−, not +2 or −3.
Tutors flagged this convention specifically. (Note the contrast with oxidation numbers, which are written the other way round: +2, −3.)
5. Electron shells
Electrons occupy shells around the nucleus, filling from the inside out.
Maximum electrons per shell:
1st shell: 2. 2nd shell: 8. 3rd shell: 8 (for the first 20 elements).
Example configurations:
- Sodium (11): 2,8,1
- Chlorine (17): 2,8,7
- Calcium (20): 2,8,8,2
For IGCSE, only write configurations for the first 20 elements — tutors noted this, and beyond calcium the pattern changes.
Why the outer shell matters:
The number of outer-shell electrons determines an element’s chemical properties and its group in the periodic table.
State the outer shell explicitly when asked. Tutors flagged that the mark often depends on naming the outermost electrons.
6. Relative atomic mass
Relative atomic mass (Ar) is the weighted average mass of an element’s atoms, taking account of the abundance of each isotope, compared with 1/12 the mass of a carbon-12 atom.
Carbon-12 is the standard against which all relative masses are measured. Uncertainty about which isotope is the standard was recorded.
The calculation:
Ar = Σ (isotope mass × % abundance) ÷ 100
Example: chlorine is 75% ³⁵Cl and 25% ³⁷Cl.
- Ar = (35 × 75 + 37 × 25) ÷ 100
- = (2625 + 925) ÷ 100 = 3550 ÷ 100 = 35.5
Divide by 100 at the end if you used percentages. Forgetting this gives an answer in the thousands.
The answer lies BETWEEN the two isotope masses, closer to the more abundant one. Here 35.5 sits between 35 and 37, nearer 35 because ³⁵Cl is more common — a good sanity check.
Estimating quickly: the Ar is close to the mass of the most abundant isotope, which tutors suggested as a way to check an answer.
Working backwards: if you know Ar and the two isotope masses, let the abundance of one be x% and the other (100 − x)%, then solve.
7. Mistakes that cost marks
Confusing proton number with nucleon number — the dominant error.
Subtracting the wrong way round when finding neutrons.
Saying isotopes have different protons.
Saying isotopes differ in electrons.
Giving only half the isotope definition.
Changing the proton number when forming an ion.
Writing an ion’s configuration before removing electrons.
Writing +2 instead of 2+ for an ion charge.
Assuming only protons contribute to mass.
Forgetting to divide by 100 in an Ar calculation.
Misreading the element from the periodic table.
Frequently asked questions
What are the three subatomic particles? Protons (+1, mass 1), neutrons (0, mass 1) and electrons (−1, negligible mass).
What is the proton number? The number of protons — the smaller number, which identifies the element.
What is the nucleon number? Protons + neutrons — the larger number.
How do I find the number of neutrons? Nucleon number − proton number.
What are isotopes? Atoms of the same element with the same protons but different neutrons.
Why do isotopes have the same chemical properties? They have the same electronic configuration.
What is an ion? An atom that has lost or gained electrons, giving it a charge.
What is the electronic configuration of Al³⁺? 2,8 — three electrons have been removed.
What is relative atomic mass? The weighted average mass of the isotopes, relative to 1/12 of a carbon-12 atom.
How do I calculate Ar? (mass × abundance) summed, divided by 100.
Quick revision checklist
- I know the charge and mass of all three particles
- I know the mass is in the nucleus, from protons and neutrons
- I can tell proton number from nucleon number
- I can calculate neutrons = A − Z
- I use the periodic table to check the element
- I can define isotopes in full, for both marks
- I know isotopes differ in neutrons only
- I can explain why their chemical properties are identical
- I know an ion changes only its electrons
- I can write an ion’s electronic configuration correctly
- I write ion charges as 2+, not +2
- I know the shell capacities 2, 8, 8
- I can calculate relative atomic mass from abundances
- I sanity-check that Ar lies between the isotope masses
These notes cover atomic structure and isotopes in the Cambridge IGCSE Chemistry (0620) syllabus and are written for Grade 9–11 / Year 10–11 students. They are based on teaching patterns observed across a large set of one-to-one IGCSE Chemistry lessons — this was the most-taught topic in that set — with particular attention to the errors students make most often and the wording examiners reward. Always check the current syllabus and data booklet for your own exam series.
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