Atomic Structure

This guide covers subtopics 2.1 to 2.3 of Topic 2, Atoms, elements and compounds, for Cambridge IGCSE Chemistry 0620 and Cambridge O Level Chemistry 5070, 2026–2028 series.

Where this fits in 0620/5070

This is the foundation the rest of Chemistry is built on — every later topic, from bonding to the mole to the Periodic Table, assumes you can describe an atom’s structure and use proton number, mass number and electronic configuration without hesitation.

Syllabus coverage

CAMBRIDGE IGCSE CHEMISTRY 0620

Core

• Describing the differences between elements, compounds and mixtures (2.1)
• The structure of the atom: a central nucleus of protons and neutrons, surrounded by electrons in shells (2.2)
• The relative charge and relative mass of a proton, neutron and electron (2.2)
• Defining proton number/atomic number and mass number/nucleon number (2.2)
• Determining electronic configuration for elements and their ions, proton number 1 to 20, e.g. 2,8,3 (2.2)
• That Group VIII noble gases have a full outer shell; outer-shell electrons equal group number for Groups I–VII; occupied shells equal period number (2.2)
• Defining isotopes, and interpreting/using isotope and ion symbols, e.g. ¹²₆C (2.3)

Subtopics 2.1 and 2.2 have no Extended-only content at all — everything above is Core, for every 0620 candidate.

Supplement / Extended

• Explaining that isotopes of the same element share chemical properties, because they have the same number of electrons and therefore the same electronic configuration (2.3)
• Calculating the relative atomic mass of an element from the relative masses and abundances of its isotopes (2.3)

CAMBRIDGE O LEVEL CHEMISTRY 5070

5070 has no Core/Extended split — all of the content above, including the two Supplement-only isotope outcomes, is required for every O Level candidate. The wording is effectively identical to the combined 0620 Core and Supplement content.

Atoms, elements, compounds and mixtures

An element contains only one type of atom. A compound forms when two or more elements are chemically bonded together in a fixed ratio, and can only be separated back into elements by a chemical reaction. A mixture contains two or more substances that are not chemically combined, in no fixed ratio, and can be separated by physical methods alone — filtration, evaporation, distillation.

Structure of the atom

An atom has a small, dense nucleus at its centre, containing protons and neutrons, surrounded by electrons arranged in shells.

Particle	Relative charge	Relative mass
Proton	+1	1
Neutron	0	1
Electron	−1	very small (often treated as negligible)

Proton number (atomic number) is the number of protons in the nucleus — this is what defines which element an atom is. Mass number (nucleon number) is the total number of protons and neutrons in the nucleus.

mass number = protons + neutrons

An atom is electrically neutral overall, so the number of electrons always equals the number of protons.

Electronic configuration

Electrons occupy shells around the nucleus, filling from the shell closest to the nucleus outward. For elements with proton number 1 to 20, each shell holds a maximum number of electrons: the first shell holds 2, the second holds 8, the third holds 8 (up to calcium).

Electronic configuration is written as a list of numbers, one per occupied shell — for example, chlorine (proton number 17) is 2,8,7: two electrons fill the first shell, eight fill the second, and the remaining seven occupy the third shell.

Three facts connect electronic configuration directly to the Periodic Table:

• Group VIII (the noble gases) have a full outer shell — this is why they’re unreactive.
• The number of outer-shell electrons equals the group number, for Groups I to VII.
• The number of occupied shells equals the period number.

Worked example. What is the electronic configuration of a potassium atom (proton number 19), and which group and period does this place it in?

19 electrons: 2 (first shell) + 8 (second shell) + 8 (third shell) = 18
remaining electrons = 19 - 18 = 1, occupying a fourth shell
configuration: 2,8,8,1

One outer-shell electron places potassium in Group I; four occupied shells place it in Period 4 — both read directly from the configuration, without needing to look either fact up separately.

Ions, not just atoms, are also required (2.2): remove electrons for a positive ion, add electrons for a negative ion, then reapply the same 2, 8, 8 filling rule. Sodium (proton number 11) is 2,8,1 as a neutral atom; the sodium ion Na⁺ has lost one electron, giving 2,8. Chlorine (17) is 2,8,7 as a neutral atom; the chloride ion Cl⁻ has gained one electron, giving 2,8,8.

Isotopes

Isotopes are atoms of the same element — same number of protons — with a different number of neutrons, and therefore a different mass number.

Isotope symbols show mass number (top) and proton number (bottom) beside the element symbol, for example ¹²₆C: carbon, proton number 6, mass number 12. The same notation extends to ions, with the charge shown as a superscript, for example ³⁵₁₇Cl⁻.

Because isotopes of an element have the same number of electrons — and therefore the same electronic configuration — they have the same chemical properties. Chemical behaviour is governed by electron arrangement, not by the number of neutrons in the nucleus.

Worked example. Chlorine exists as two isotopes: ³⁵Cl (75% abundance) and ³⁷Cl (25% abundance). Calculate its relative atomic mass.

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

This calculated value, 35.5, is exactly the Ar you’ll find quoted for chlorine in a data book — relative atomic mass is a weighted average across an element’s naturally occurring isotopes, which is why it isn’t always a whole number.

Common mistakes

• Confusing proton number and mass number. Proton number identifies the element and never changes for that element; mass number varies between isotopes of the same element.
• Confusing mass number with relative atomic mass. Mass number is a whole number belonging to one specific atom; relative atomic mass is a weighted average across an element’s naturally occurring isotopes, so it isn’t always a whole number.
• Forgetting electrons have negligible mass but real (negative) charge. Both facts matter — mass number ignores electrons, but overall charge does not.
• Forgetting to adjust the electron count when writing an ion’s configuration. Positive ions have lost electrons; negative ions have gained them — the proton count (and therefore the element) never changes.
• Writing an electronic configuration that exceeds shell capacity. The first shell never holds more than 2, and (up to calcium) the second and third shells never hold more than 8 each.
• Assuming isotopes have different chemical properties because their mass differs. Chemical behaviour depends on electron arrangement, which is identical across an element’s isotopes.
• Averaging isotope masses without weighting by abundance. A simple average of 35 and 37 gives 36, not chlorine’s actual Ar of 35.5 — abundance must be factored in.

Quick revision checklist

• Element, compound, mixture — and how mixtures are separated
• Proton, neutron, electron: relative charge and mass
• Proton number and mass number, defined and distinguished
• Electronic configuration by shell, for proton numbers 1–20, including ions
• Reading group number and period number from electronic configuration
• Isotopes: definition, symbol notation, and why chemical properties don’t change between them
• (0620 Extended, 5070 required) calculating relative atomic mass from isotopic abundances

What you should know before moving to AS Chemistry

By the end of this topic you should be comfortable writing an electronic configuration by shell (e.g. 2,8,7) for any element up to proton number 20, explaining what proton number and mass number mean, and calculating relative atomic mass from isotopic abundance data. AS Chemistry takes electronic configuration much further — into sub-shells and orbitals — and uses ionisation energy data in ways this topic doesn’t touch at all.

Continue to AS Chemistry

If you’re moving on to Cambridge International AS & A Level Chemistry 9701, Atomic Structure: Particles, Radius and Isotopes picks up directly from here — the same particles and isotopes, now with electric-field behaviour and atomic/ionic radius trends explained rather than just described. Atomic Structure: Orbitals and Ionisation Energy follows it, replacing the shell-only picture above with shells, sub-shells and orbitals, and introducing ionisation energy as measurable data you’re expected to interpret. Neither is required for IGCSE or O Level; they’re there for when you need them.

Related resources

• Formulae, Equations and the Mole — Ar and isotopes applied to reacting masses and the mole
• Ionic, Covalent and Metallic Bonding — electronic configuration applied to bonding (subtopics 2.4–2.7)
• The Periodic Table: Groups and Trends — Group I, Group VII and Group VIII in depth
• Atomic Structure (OCR GCSE) — the equivalent subtopic from a different awarding body
• Cambridge IGCSE Chemistry hub · Cambridge O Level Chemistry hub

Written against Cambridge IGCSE Chemistry 0620 and Cambridge O Level Chemistry 5070, 2026–2028 series. Always check the current syllabus for your examination year.