Atomic Structure and the Development of the Atom
The Structure of the Atom
An atom is the smallest part of an element that can exist. Atoms are extraordinarily small — a typical atomic radius is about \(1 \times 10^{-10}\) m. Each atom has a tiny, dense nucleus at its centre. The nucleus contains two types of sub-atomic particle: protons (positive charge, relative mass 1) and neutrons (no charge, relative mass 1). Electrons occupy energy levels around the nucleus and have a negative charge. Their relative mass is so small — roughly one part in 1836 that of a proton — that it is treated as negligible.
Relative charges and approximate relative masses: Proton — charge +1, mass 1. Neutron — charge 0, mass 1. Electron — charge −1, mass negligible. The nucleus is much smaller than the atom as a whole; most of the atom is empty space.
Atomic Number and Mass Number
The atomic number (proton number) tells you how many protons are in the nucleus. It defines what element an atom is — all atoms of the same element have the same atomic number. In a neutral atom, the number of electrons equals the number of protons, so there is no overall charge. The mass number is the total number of protons plus neutrons. From these two values you can work out the number of neutrons:
$$\text{number of neutrons} = \text{mass number} - \text{atomic number}$$
For example, a sodium atom has atomic number 11 and mass number 23, giving it 12 neutrons, 11 protons and 11 electrons. Standard notation places the mass number at the top left and the atomic number at the bottom left of the symbol: \(^{23}_{11}\text{Na}\).
Isotopes
Isotopes are atoms of the same element that have the same atomic number but different mass numbers because they contain different numbers of neutrons. For example, carbon-12 (\(^{12}_{6}\text{C}\)) has 6 neutrons while carbon-14 (\(^{14}_{6}\text{C}\)) has 8 neutrons. Because isotopes have the same number of protons and electrons, they have identical chemical properties — they react in the same way. Their physical properties (such as density and diffusion rate) differ slightly because of the mass difference. Radioactive isotopes have unstable nuclei that decay, emitting radiation.
Isotopes of the same element have the same number of protons (same atomic number) but different numbers of neutrons (different mass numbers). They have identical chemical properties because they have the same number and arrangement of electrons.
The Development of the Atomic Model
Our understanding of the atom developed gradually as new experimental evidence challenged existing ideas. This is a classic example of how science works — models are revised or replaced when they cannot explain new observations.
Dalton’s Model
John Dalton proposed that atoms were tiny, solid, indivisible spheres. He thought all atoms of the same element were identical and that atoms of different elements had different masses.
Thomson’s Plum Pudding Model
J.J. Thomson discovered the electron and proposed a model in which negative electrons were scattered throughout a uniform sphere of positive charge, like plums in a pudding.
Rutherford’s Nuclear Model
Ernest Rutherford, Hans Geiger and Ernest Marsden fired positively charged alpha particles at a thin sheet of gold foil. Most passed straight through (confirming the atom is mostly empty space), but some deflected at large angles and a small number bounced directly back. This disproved the plum pudding model and led Rutherford to propose a new model: a tiny, dense, positively charged nucleus at the centre with electrons orbiting around it in a mostly empty atom.
Bohr’s Model
Niels Bohr refined Rutherford’s model by placing electrons in specific fixed energy levels (shells) at defined distances from the nucleus. This explained why atoms emit light at particular wavelengths.
Chadwick’s Neutron
James Chadwick provided experimental evidence for the existence of the neutron, completing the modern nuclear model with protons and neutrons in the nucleus.
The alpha scattering experiment (sometimes called the Geiger–Marsden experiment) is central to this topic. The key observations were: (1) most alpha particles passed straight through — the atom is mostly empty space; (2) some deflected at large angles — they came close to a concentrated positive charge; (3) a tiny fraction bounced almost straight back — they hit something very small, dense and positive head-on. If the plum pudding model were correct, the positive charge would be spread out and alpha particles would be barely deflected. The large deflections could only be explained by a small, dense, positively charged nucleus. This is a key example of how a scientific model is overturned by experimental evidence that it cannot explain.
Electronic Configuration
Electrons occupy energy levels (shells) around the nucleus. The first shell can hold up to 2 electrons; the second and third shells can each hold up to 8 electrons. Electrons fill the lowest available energy level first. You write an atom’s electronic configuration as a sequence of numbers showing how many electrons are in each shell. For example, sodium (11 electrons) is written 2, 8, 1.
Start with the atomic number (which equals the number of electrons in a neutral atom). Fill shell 1 with up to 2 electrons, then shell 2 with up to 8, then shell 3 with up to 8. Examples: oxygen (8 electrons) → 2, 6. Sodium (11 electrons) → 2, 8, 1. Chlorine (17 electrons) → 2, 8, 7. Calcium (20 electrons) → 2, 8, 8, 2. The number of electrons in the outer shell matches the group number for Groups 1 to 7, which explains why elements in the same group have similar chemical properties — they react in the same way because they have the same number of outer electrons.
The electronic configuration of an element determines how it bonds. Group 1 elements (one outer electron) tend to lose that electron and form +1 ions. Group 7 elements (seven outer electrons) tend to gain one electron to form −1 ions. Group 0 elements have complete outer shells, which is why the noble gases are very unreactive. A good working knowledge of electronic configuration lets you predict whether an element will form ionic or covalent bonds and what charges its ions will carry.