Atomic Structure and Models Through Time
The Structure of the Atom
An atom is the smallest particle of an element that retains the chemical properties of that element. Atoms are extraordinarily small — a typical atomic radius is around \(1 \times 10^{-10}\) m. At the centre of each atom is a tiny, dense nucleus. The nucleus contains two types of subatomic particle: protons, which carry a positive charge, and neutrons, which carry no charge. Electrons orbit the nucleus in energy levels (also called shells) and carry a negative charge. The nucleus accounts for almost all of the mass of an atom, yet it occupies only a tiny fraction of the atom’s volume — the rest is essentially empty space.
Subatomic particles: Proton — charge +1, relative mass 1. Neutron — charge 0, relative mass 1. Electron — charge −1, relative mass approximately \(\frac{1}{1836}\) (negligible). The nuclear radius is far smaller than the atomic radius — most of an atom is empty space.
Atomic Number, Mass Number and Neutrons
The atomic number (also called the proton number) tells you how many protons are in the nucleus. It uniquely identifies the element — all atoms of the same element have the same atomic number. In a neutral atom, the number of protons equals the number of electrons, so the atom has no overall charge. The mass number is the total count of protons plus neutrons in the nucleus. To find the number of neutrons:
$$\text{neutrons} = \text{mass number} - \text{atomic number}$$
For example, a chlorine atom with mass number 35 and atomic number 17 has \(35 - 17 = 18\) neutrons, 17 protons, and 17 electrons. Standard notation places the mass number top-left and the atomic number bottom-left of the symbol: \(^{35}_{17}\text{Cl}\).
Isotopes
Isotopes are atoms of the same element with different numbers of neutrons. Because the number of protons (and therefore electrons) is unchanged, isotopes have identical chemical properties. Only their mass differs. For example, carbon-12 has 6 protons and 6 neutrons, while carbon-14 has 6 protons and 8 neutrons. Both are carbon; both react in the same way chemically. Chlorine exists in nature as two isotopes — chlorine-35 and chlorine-37 — which is why the relative atomic mass of chlorine given in the periodic table is approximately 35.5 (a weighted average of the two). Isotopes are important in applications ranging from medical imaging to carbon dating.
An ion is an atom (or group of atoms) that has gained or lost electrons and therefore carries an overall electric charge. Losing electrons produces a positive ion (cation) because there are now more protons than electrons. Gaining electrons produces a negative ion (anion). The nucleus does not change when an ion forms — only the electron count changes. For example, sodium loses one electron to form \(\text{Na}^+\) and chlorine gains one electron to form \(\text{Cl}^-\). The charges are shown as superscripts: \(\text{Mg}^{2+}\) has lost two electrons; \(\text{O}^{2-}\) has gained two.
How the Atomic Model Developed
Our current understanding of the atom was not established in a single discovery. It developed gradually over more than a century as new experimental evidence challenged existing ideas and forced scientists to revise their models. This process illustrates a key principle of science: a model is only useful as long as it fits the available evidence. When new evidence arrives that a model cannot explain, the model must change. This is not a failure — it is how science makes progress.
Dalton’s Model
John Dalton proposed that atoms were tiny, solid, indivisible spheres. He argued that all atoms of the same element were identical and that atoms of different elements had different masses. This model explained why elements always combine in fixed ratios, but gave no internal structure for the atom.
Thomson’s Plum Pudding Model
J.J. Thomson discovered the electron by experimenting with cathode rays. He realised atoms contained smaller, negatively charged particles and proposed a new model: a ball of diffuse positive charge with negative electrons embedded in it — like plums in a pudding. This was the first model to include subatomic particles.
Rutherford’s Nuclear Model
Ernest Rutherford, working with Hans Geiger and Ernest Marsden, fired positively charged alpha particles at a thin gold foil. The result was unexpected: most particles passed straight through, but some were deflected at large angles and a small number bounced almost directly back. The plum pudding model predicted that alpha particles would pass through the diffuse positive charge with only minor deflections. The actual results could only be explained if the positive charge and most of the mass were concentrated in a tiny, dense region — the nucleus. Rutherford’s nuclear model replaced the plum pudding model.
Bohr’s Model
Niels Bohr refined the nuclear model by showing that electrons could not orbit at any distance — they were confined to specific energy levels (fixed shells) at set distances from the nucleus. Electrons in a given shell have a fixed energy; they can absorb energy and jump to a higher shell, or fall back and release energy as light. Bohr’s model explained the emission spectra of elements, which the Rutherford model could not.
The alpha particle scattering experiment showed that most alpha particles passed straight through the gold foil (atom is mostly empty space) but some bounced back sharply (small, dense, positive nucleus). This disproved Thomson’s plum pudding model and led to the nuclear model.
Models as Tools: Strengths and Limitations
Each atomic model was the best explanation of the evidence available at the time. Dalton’s model was useful for explaining chemical reactions but could not account for electrical phenomena. Thomson’s model incorporated electrons but placed them in the wrong configuration. Rutherford’s model identified the nucleus but could not explain why electrons did not spiral into it. Bohr’s model resolved this by proposing fixed energy levels. Later developments in quantum mechanics produced an even more sophisticated picture — but at GCSE, the Bohr model is the accepted working model. Understanding that models change with evidence is as important as knowing the models themselves.
Electronic Configuration
Electrons occupy shells around the nucleus. The first shell holds a maximum of 2 electrons; the second and third shells each hold a maximum of 8. Electrons always fill the lowest available shell first. The electronic configuration of an element determines how it reacts chemically.
Use the atomic number to find the total number of electrons in a neutral atom. Fill shells in order: first shell up to 2, then second up to 8, then third up to 8. Examples: sodium (atomic number 11) → 2, 8, 1. Calcium (atomic number 20) → 2, 8, 8, 2. The number of electrons in the outer shell equals the group number for Groups 1–7. This is why elements in the same group have similar chemical behaviour: they have the same number of outer-shell electrons available for bonding.