Atomic Structure
Welcome to Atomic Structure & The Periodic Table — the foundation of all chemistry. Everything you touch, breathe, eat and drink is made of atoms, and understanding how atoms are built is the key to unlocking every other topic in this subject.
In this module you will learn how atoms are structured with protons, neutrons and electrons, how the periodic table organises elements by their properties, how electron configurations determine chemical behaviour, and how isotopes and relative atomic mass calculations work. These ideas run through the whole chemistry syllabus.
By the end you will be comfortable drawing electron shell diagrams, predicting element properties from their position in the periodic table, performing mass-number and mole calculations, and explaining ionic, covalent and metallic bonding and the shapes of simple molecules with confidence. Let's get started!
What the Exam Tests
The official chemistry syllabus has four modules: Basic Chemical Concepts and Calculations, Properties and Reactions of Substances, Chemical Theories and Laws, and Chemical Experiments and Applications. This course module teaches the start of Chemical Theories and Laws: atomic structure and the periodic law, and chemical bonds and intermolecular forces. Its Classifying Matter lesson and its Relative Formula Mass & Moles lesson belong to Basic Concepts and Calculations (the classification and changes of state of matter; calculations involving the amount of substance).
Every question is single-answer multiple choice, set in Chinese or English. Questions can ask you to count the protons, neutrons and electrons in an ion, write an electron configuration, compare atomic radius or electronegativity across the periodic table, decide whether a bond is ionic or covalent, or predict the shape of a molecule.
In our mock papers Chemical Theories and Laws gets 14 of the 48 questions and Basic Concepts and Calculations 12. The CSCA publishes the topics but not how many questions each gets, so this split is our own reading of the syllabus.
An atom is the smallest particle of an element that retains the chemical properties of that element. Every atom consists of a tiny, dense nucleus at the centre, surrounded by electrons that orbit in shells (energy levels).
Sub-atomic Particles
- Proton (p⁺) — found in the nucleus, relative mass ≈ 1, relative charge = +1.
- Neutron (n⁰) — found in the nucleus, relative mass ≈ 1, relative charge = 0. Not every nucleus has one: the commonest hydrogen atom, ¹₁H (protium, 氕), is a single proton with no neutron.
- Electron (e⁻) — found in shells around the nucleus, relative mass ≈ 1/1836 (negligible), relative charge = −1.
Key Numbers
- Atomic number (Z) = number of protons. This defines which element the atom is. All carbon atoms have Z = 6.
- Mass number (A) = protons + neutrons. This tells you the total "heavy" particles in the nucleus.
- Number of neutrons (N) = A − Z.
In Chinese textbooks: 原子序数 = 核电荷数 = 质子数 = 核外电子数 (for a neutral atom), and 质量数 (A) = 质子数 (Z) + 中子数 (N).
In a neutral atom, the number of electrons equals the number of protons (charges balance). When an atom gains or loses electrons it becomes an ion — a cation (+) if it loses electrons, an anion (−) if it gains them.
Notation
We write an element as AZX where A is the mass number (top) and Z is the atomic number (bottom). For example, carbon-12 is written ¹²₆C — it has 6 protons, 6 neutrons and (when neutral) 6 electrons.
The nucleus is incredibly small compared to the atom. If the atom were the size of a football stadium, the nucleus would be a marble on the centre spot — yet it contains over 99.9% of the atom's mass.
Understanding this basic structure is essential because the number and arrangement of electrons determines how atoms bond, react and form compounds.
How the model of the atom changed
- Dalton (1803): atoms are solid, indivisible spheres, and all atoms of an element are identical. Isotopes later showed the second half of that is not quite true.
- Thomson (1897): discovered the electron, so the atom could not be indivisible. His "plum pudding" model had electrons embedded in a ball of positive charge.
- Rutherford (1911): the gold foil experiment fired alpha particles at thin gold. Most passed straight through, but a very few bounced almost straight back — impossible if the positive charge were spread out. The conclusion was the nuclear model: a tiny, dense, positively charged nucleus with the atom mostly empty space.
- Bohr (1913): electrons orbit in fixed energy levels, or shells, and can only move between them by absorbing or emitting a photon of exactly the right energy. This is what makes atomic spectra sets of sharp lines rather than a continuous band.
- The modern picture: electrons occupy an electron cloud of orbitals — regions where an electron is likely to be found — rather than travelling in neat circles.
A typical question of this kind gives you an observation and asks which model it supports, so learn each model by the evidence that killed the one before it.
Key takeaway: Atoms have protons (+1) and neutrons (0) in the nucleus, electrons (−1) in shells. Atomic number = protons; mass number = protons + neutrons.
Key Formulas
Z = atomic number = protons | A = mass number = protons + neutrons | Neutrons = A − Z | In neutral atom: electrons = protons
Worked Example 1
Example 1: An atom of sodium has atomic number 11 and mass number 23. How many protons, neutrons and electrons does it have?
Protons = Z = 11. Neutrons = A − Z = 23 − 11 = 12. Electrons = 11 (neutral atom, so electrons = protons).
Worked Example 2
Example 2: A chloride ion Cl⁻ has Z = 17, A = 35. How many electrons does it have?
Neutral Cl has 17 electrons. The ⁻ charge means it gained 1 electron → 18 electrons.
Worked Example 3
Example 3: An atom has 26 protons and 30 neutrons. Identify the element and its mass number.
Z = 26 → Iron (Fe). A = 26 + 30 = 56. This is ⁵⁶₂₆Fe.
Exam tip: Always check whether the question says "atom" or "ion" — ions have a different electron count.
Exam tip: Mass number is always a whole number. Relative atomic mass (on the periodic table) is usually a decimal because it is a weighted average of isotopes.
Common mistake: Confusing mass number with relative atomic mass: Mass number is for one specific atom (always a whole number). Relative atomic mass is the weighted average across all isotopes.
Common mistake: Forgetting that ions change electron count, not proton count: A Na⁺ ion still has 11 protons; it just lost an electron to have only 10.
Try This Yourself
An atom of potassium (K) has Z = 19 and A = 39. Calculate the number of protons, neutrons and electrons.
Module Summary
Open and read all sections to complete this module