What you'll learn
Atomic structure and the periodic table form the foundation of GCSE Chemistry: almost every other topic depends on understanding what atoms are made of and how the periodic table is organised. In this guide you will learn the structure of the atom, the properties of protons, neutrons and electrons, how to work out electronic structures, how the modern periodic table is arranged, how our model of the atom developed historically, and the meaning of atomic number, mass number and isotopes. Mastering these ideas gives you the language and concepts needed for bonding, reactions and the rest of the course.
Key terms and definitions
Atom — the smallest part of an element that can exist.
Proton — a positively charged subatomic particle found in the nucleus (relative mass 1).
Neutron — a neutral subatomic particle found in the nucleus (relative mass 1).
Electron — a negatively charged subatomic particle that occupies shells around the nucleus (relative mass very small, ~1/1835).
Atomic number — the number of protons in an atom (defines the element).
Mass number — the total number of protons and neutrons.
Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons.
Relative atomic mass (Aᵣ) — the average mass of atoms of an element, taking isotopic abundance into account.
Core concepts
The structure of the atom
Atoms have a tiny central nucleus containing protons and neutrons, surrounded by electrons arranged in shells (energy levels). The nucleus is positively charged and contains almost all the mass, but is extremely small compared with the whole atom — the radius of an atom is about 0.1 nanometres (1 × 10⁻¹⁰ m), while the nucleus is around 1 × 10⁻¹⁴ m, roughly 1/10000 of the atom's radius. Atoms have no overall charge because the number of protons equals the number of electrons.
Atomic number, mass number and isotopes
The atomic number is the number of protons and identifies the element. The mass number is protons plus neutrons, so the number of neutrons = mass number − atomic number. Isotopes are atoms of the same element with different numbers of neutrons; they have identical chemical properties because they have the same electronic structure, but slightly different masses. The relative atomic mass is a weighted mean of the isotope masses according to their abundance.
Electronic structure
Electrons occupy shells, filling the lowest energy level first. The first shell holds up to 2 electrons, the second up to 8, and the third up to 8 (at GCSE). For example, sodium (atomic number 11) has the configuration 2,8,1. The number of electrons in the outer shell determines an element's chemical behaviour and its group in the periodic table.
The development of the atomic model
Ideas about the atom changed as new evidence emerged. John Dalton described atoms as solid spheres. J.J. Thomson discovered the electron and proposed the "plum pudding" model. The alpha scattering experiment (Rutherford, Geiger and Marsden) showed most of the atom is empty space with a small dense positive nucleus, replacing the plum pudding model. Niels Bohr proposed electrons orbit at fixed energy levels, and later experiments established the proton; James Chadwick discovered the neutron.
Arrangement of the periodic table
Elements are arranged in order of increasing atomic number. Vertical columns are groups; elements in the same group have the same number of outer-shell electrons and similar chemical properties. Horizontal rows are periods. Metals are on the left and centre; non-metals on the right. Mendeleev produced an early periodic table, ordering elements by atomic mass but leaving gaps for undiscovered elements and even predicting their properties — his arrangement was vindicated when those elements were found.
Worked examples
Example 1: Working out neutrons
An atom has atomic number 17 and mass number 35. How many protons, neutrons and electrons does it have?
Protons = atomic number = 17. Electrons = protons = 17. Neutrons = mass number − atomic number = 35 − 17 = 18.
Example 2: Electronic structure
Write the electronic structure of an atom with 19 electrons.
Fill shells in order: 2, then 8, then 8, leaving 1: 2,8,8,1. This element (potassium) is in Group 1 because it has one outer electron.
Example 3: Relative atomic mass
Chlorine is 75% ³⁵Cl and 25% ³⁷Cl. Calculate its relative atomic mass.
Aᵣ = (75 × 35 + 25 × 37) ÷ 100 = (2625 + 925) ÷ 100 = 3550 ÷ 100 = 35.5.
Common mistakes and how to avoid them
Confusing atomic number and mass number. Atomic number = protons; mass number = protons + neutrons. Neutrons are the difference.
Overfilling shells. Remember 2, then 8, then 8 at GCSE; don't put more than the limit in a shell.
Thinking isotopes behave differently chemically. They react the same because chemistry depends on electrons, which are identical; only mass differs.
Muddling the historical models. Link each scientist to their contribution and the evidence that prompted the change.
Forgetting weighting in relative atomic mass. Multiply each isotope mass by its abundance, not a simple average.
Exam technique for Atomic Structure and the Periodic Table
State particle properties precisely — relative charges (+1, 0, −1) and relative masses (1, 1, very small).
Show electronic structures clearly, as numbers (2,8,1) or shell diagrams, and link the outer electrons to group number.
Quote evidence for the models, especially how alpha scattering disproved the plum pudding model.
Calculate relative atomic mass using abundance weighting, showing your working.
Use correct vocabulary — group, period, isotope, nucleus — to access the higher marks.
Quick revision summary
An atom has a tiny, dense, positively charged nucleus of protons (+1, mass 1) and neutrons (0, mass 1), surrounded by electrons (−1, tiny mass) in shells. The atomic number (protons) defines the element; the mass number is protons + neutrons; neutrons = mass number − atomic number. Isotopes share proton number but differ in neutrons, behaving identically chemically; relative atomic mass is the abundance-weighted average of isotope masses. Electrons fill shells 2, 8, 8, and the outer-shell count sets the group. The atomic model evolved from Dalton's spheres, through Thomson's plum pudding, to the nuclear atom revealed by alpha scattering, refined by Bohr's energy levels and Chadwick's neutron. The periodic table orders elements by atomic number into groups (same outer electrons, similar properties) and periods, with metals to the left and non-metals to the right; Mendeleev anticipated it by leaving gaps for undiscovered elements. Know particle properties, write electronic structures, link them to position in the table, recall the evidence behind each model, and calculate relative atomic mass with proper weighting.