Atomic Structure

Summary: Atoms consist of protons, neutrons, and electrons. Atomic number (Z) equals protons; mass number equals protons + neutrons. Electrons are arranged in shells (2.8.8…). Isotopes are atoms of the same element with different numbers of neutrons. Tags: igcse chemistry atomic-structure Created: 2026-07-14 Last Updated: 2026-07-18

Atomic structure describes how an atom — the smallest particle of an element that retains its chemical identity — is built from a tiny, dense, positively charged nucleus surrounded by a cloud of negatively charged electrons. The nucleus contains two types of nucleon: the positively charged Proton and the uncharged Neutron, which together account for almost all of the atom’s mass. The number of protons defines the Atomic Number (Proton Number) (Z), which identifies the element, while the total count of protons and neutrons gives the Mass Number (Nucleon Number) (A); both are displayed compactly using Nuclide Notation. Electrons occupy a series of electron shells (energy levels) around the nucleus, filling in a fixed pattern (2, 8, 8, …) that is summarised by an atom’s Electronic Configuration and that determines its position in the Periodic Table and its chemical behaviour. Atoms of the same element can differ in neutron count — these variants are called isotopes — and the weighted average of their masses gives the element’s Relative Atomic Mass (Ar). Understanding this structure is the foundation of IGCSE Chemistry, explaining everything from bonding and ion formation to periodic trends and why Ar values are often not whole numbers.

Subatomic Particles

All matter is made of atoms. Atoms themselves are composed of three subatomic particles:

Key points:

• Most of the mass of an atom is concentrated in the nucleus (protons + neutrons)
• Most of the volume of an atom is empty space — the electrons orbit far from the nucleus
• The nucleus is tiny compared to the size of the atom: if an atom were the size of a football stadium, the nucleus would be about the size of a pea
• Atoms are electrically neutral overall: number of protons = number of electrons

Atomic Number and Mass Number

Every element is defined by its atomic number (Z):

• Atomic Number (Z) = number of protons in the nucleus
• In a neutral atom, number of protons = number of electrons
• Z determines which element it is — change the number of protons and you change the element

The mass number (A) (also called nucleon number) tells you the total number of particles in the nucleus:

• Mass Number (A) = number of protons + number of neutrons
• Number of neutrons = A − Z

Nuclude notation (also called atomic notation):

  Mass Number → A
                 X   ← Element symbol
  Atomic Number → Z

Example: ²³₁₁Na

• Z = 11 → 11 protons, 11 electrons (in a neutral atom)
• A = 23 → total nucleons
• Neutrons = 23 − 11 = 12 neutrons

Example: An atom has 26 protons, 30 neutrons, and 26 electrons.

• (a) What is its atomic number? → 26
• (b) What is its mass number? → 26 + 30 = 56
• (c) What element is this? → Iron (Fe)

Example: An atom of chlorine has 17 protons and 20 neutrons.

• (a) State its atomic number: 17
• (b) State its mass number: 17 + 20 = 37
• (c) Write its nuclide notation: ³⁷₁₇Cl

Electron Configuration

Electrons are arranged in shells (also called energy levels) around the nucleus.

The rules for filling electron shells (for the first 20 elements):

1. 1st shell: holds a maximum of 2 electrons
2. 2nd shell: holds a maximum of 8 electrons
3. 3rd shell: holds a maximum of 8 electrons (for the first 20 elements)

How to write electron configuration:

• Count the total number of electrons (equals atomic number for a neutral atom)
• Fill shells in order: 1st shell first, then 2nd, then 3rd
• Write as: electron configuration = 2.8.X (or whatever the numbers are)

Examples of electron configurations for the first 20 elements:

What electron configuration tells you:

• Number of outer-shell electrons = Group number (for Groups 1–2 and 13–18)
  • Group 1: 1 outer electron (e.g., Na: 2.8.1)
  • Group 7: 7 outer electrons (e.g., Cl: 2.8.7)
  • Group 0 (Noble gases): full outer shell — 2 or 8 electrons (e.g., Ne: 2.8, Ar: 2.8.8)
• Number of occupied shells = Period number
  • Period 1: 1 shell (H, He)
  • Period 2: 2 shells (Li through Ne)
  • Period 3: 3 shells (Na through Ar)

Example: An element has electron configuration 2.8.6.

• (a) How many electrons does it have? → 2 + 8 + 6 = 16 electrons
• (b) Which element is it? → Sulfur (S), Z = 16
• (c) What group is it in? → Group 6 (Group 16) — 6 electrons in the outer shell
• (d) What period is it in? → Period 3 — 3 occupied shells

Isotopes

Definition: Isotopes are atoms of the same element (same number of protons / same atomic number) that have different numbers of neutrons (different mass numbers).

Examples of isotopes:

Properties of isotopes:

• Same chemical properties — because they have the same number of electrons (specifically, the same number of outer electrons), and chemical reactions involve electrons
• Different physical properties — because they have different masses (different number of neutrons), so properties like density, rate of diffusion, and mass spectra differ

Notable isotopes:

• Carbon-14 dating: ¹⁴C is radioactive (half-life ~5730 years). Used to estimate the age of archaeological artefacts containing once-living material
• Iodine-131: Radioactive isotope used in medical imaging and thyroid treatment
• Chlorine-35 and chlorine-37: Reason why the relative atomic mass of chlorine is 35.5 — it’s a weighted average

Relative Atomic Mass (Ar)

The relative atomic mass (Ar) is the weighted average mass of all the isotopes of an element, compared to 1/12 the mass of a carbon-12 atom.

Formula:

Ar = (mass × abundance of isotope 1) + (mass × abundance of isotope 2) + ...
     ---------------------------------------------------------------
                        total abundance (100%)

Example: Chlorine has two isotopes: ³⁵Cl (75%) and ³⁷Cl (25%). Calculate the relative atomic mass.

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

This is why many Ar values on the Periodic Table are not whole numbers — they are weighted averages of isotopes.

Key Concepts from Past Papers

• Atomic number: the number of protons in the nucleus of an atom
• Mass number: the total number of protons and neutrons in the nucleus
• Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons
• Relative atomic mass: the weighted average mass of isotopes relative to 1/12 the mass of carbon-12

Keywords from Past Papers

group, protons, electrons, hydrogen, carbon, cooh, down, metal, circle, around, mixture, neutrons, properties, elements, similar

Related Notes

• Atom (Chemistry) — The smallest particle of an element
• Nucleus (Chemistry) — The dense centre of the atom containing protons and neutrons
• Proton — Positively charged nucleon; count defines the element
• Neutron — Uncharged nucleon; varies between isotopes
• Electron — Negatively charged particle in shells around the nucleus
• Electron Shell — Energy levels holding 2, 8, 8… electrons
• Atomic Number (Proton Number) — Z, the number of protons
• Mass Number (Nucleon Number) — A, protons + neutrons
• Electronic Configuration — Shell-by-shell arrangement of electrons (e.g., 2.8.1)
• Isotope — Same element, different number of neutrons
• Nuclide Notation — Symbol form showing mass number and atomic number (e.g., ³⁷₁₇Cl)
• Relative Atomic Mass — Weighted average mass of an element’s isotopes
• Ions and Ionic Bonds — How atoms gain/lose electrons to form ions
• Covalent Bonding — How atoms share electrons
• Metallic Bonding — Delocalised electrons in metal lattices
• Giant Structures — Diamond, graphite, silicon dioxide
• Periodic Table Trends — How electron configuration determines periodic trends
• IGCSE-Chem-Index — Full IGCSE Chemistry index

Sources

• OpenStax Chemistry 2e — Chapter 2: Atoms, Molecules, and Ions, Rice University (free, CC BY 4.0)
• BBC Bitesize GCSE Chemistry — Atomic Structure, BBC (free educational resource)
• Cambridge IGCSE Chemistry 0620 — Syllabus Section 2.2-2.3 (Atoms, elements and compounds), Cambridge Assessment International Education
• CK-12 Chemistry for High School — Atomic Structure, CK-12 Foundation (free, CC BY-NC 3.0)

Past Paper Sources

• 0620/31 May/June 2015: Q22(b)(iii) (0m), Q33(b)(ii) (0m), Q44(a)(i) (0m) (+2 more)
• 0620/32 Feb/March 2015: Q11(b)(ii) (3m), Q66(b)(i) (1m)
• 0620/32 Feb/March 2017: Q33(b)(i) (1m), Q77(c)(i) (1m)
• 0620/32 Feb/March 2018: Q22(b)(i) (1m), Q33(b)(ii) (1m), Q55(d)(i) (1m)
• 0620/32 Feb/March 2019: Q22(a)(i) (2m)
• 0620/32 Feb/March 2020: Q44(a)(i) (1m)
• 0620/32 Feb/March 2021: Q33(b)(ii) (1m), Q66(a)(i) (1m)
• 0620/32 Feb/March 2022: Q33(b)(i) (1m)
• 0620/32 Feb/March 2023: Q66(a)(i) (1m)
• 0620/32 May/June 2018: Q22(c)(ii) (2m), Q33(a)(ii) (1m), Q77(a)(iii) (1m) (+1 more)
• 0620/32 May/June 2020: Q44(a)(i) (1m)
• 0620/33 May/June 2016: Q11(b)(i) (1m), Q33(a)(ii) (1m)

Common Misconceptions