Atoms, Elements and the Development of the Atomic Model
Atoms, Elements and Compounds
All matter is made of tiny particles called atoms. An element is a pure substance containing only one type of atom — for example, copper contains only copper atoms and chlorine contains only chlorine atoms. When atoms of different elements join together chemically, they form a compound. In a compound, the atoms are held by chemical bonds and can only be separated by chemical reactions. Mixtures, by contrast, contain more than one substance not chemically bonded, and can be separated by physical means.
All matter is made of atoms. Elements contain only one type of atom. Compounds contain atoms of two or more different elements chemically bonded together in fixed proportions. Atoms are neither created nor destroyed in chemical reactions.
Structure of the Atom
The atom has a central nucleus containing protons (positively charged) and neutrons (no charge). Electrons (negatively charged) orbit the nucleus in shells (energy levels). In a neutral atom, the number of protons equals the number of electrons, so the charges balance.
The relative masses and charges of the three subatomic particles are:
- Proton: relative mass 1, charge +1.
- Neutron: relative mass 1, charge 0.
- Electron: relative mass approximately 0 (very small compared to protons and neutrons), charge −1.
Two numbers describe every atom. The atomic number (proton number) is the number of protons in the nucleus and determines which element the atom is. The mass number is the total number of protons plus neutrons. The number of neutrons in a neutral atom = mass number − atomic number.
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. For example, carbon-12 has 6 protons and 6 neutrons, while carbon-14 has 6 protons and 8 neutrons. Both are carbon because they have atomic number 6. Isotopes of an element have the same chemical properties (same electron arrangement) but different physical properties (different masses).
Because naturally occurring elements often consist of a mixture of isotopes, the relative atomic mass (Aᵣ) is calculated as a weighted mean. For example, chlorine has two isotopes: chlorine-35 (about 75% abundant) and chlorine-37 (about 25% abundant). Its relative atomic mass is approximately: $$A_r = \frac{(75 \times 35) + (25 \times 37)}{100} = \frac{2625 + 925}{100} = 35.5$$
Isotopes have the same proton number (atomic number) but different neutron numbers (different mass numbers). Relative atomic mass is the weighted mean mass, taking account of the abundance of each isotope. Isotopes have identical chemistry because they have the same electron arrangement.
The Development of the Atomic Model
Our model of the atom has changed repeatedly as new experimental evidence came to light. This history shows how scientific models are provisional — they are the best explanation of available evidence and are revised when new evidence demands it.
Dalton’s Atomic Theory
John Dalton proposed that all matter was made of tiny indivisible spheres called atoms. A key claim of Dalton’s model was that all atoms of a given element are identical in mass and properties, and that atoms of different elements differ in mass. He also proposed that atoms cannot be created or destroyed in chemical reactions and that compounds are formed by the combination of atoms in fixed whole-number ratios. The existence of isotopes later showed that not all atoms of an element are identical in mass — refining Dalton’s claim.
Thomson — Discovery of the Electron and the ‘Plum Pudding’ Model
J.J. Thomson discovered the electron and concluded that atoms contained smaller, negatively charged particles. He proposed the ‘plum pudding’ model: the atom is a diffuse sphere of positive charge with electrons embedded in it like plums in a pudding. This model retained the idea that the atom is a solid structure but revealed that atoms have internal structure.
Rutherford — the Gold Foil Experiment
Ernest Rutherford, working with Hans Geiger and Ernest Marsden, fired alpha particles at a thin sheet of gold foil. Most particles passed straight through, but a small number were deflected at large angles, and a very few bounced almost straight back. This was inconsistent with the plum pudding model. Rutherford concluded that most of the mass of an atom is concentrated in a tiny, dense, positively charged nucleus, with a large amount of empty space around it. Electrons orbit the nucleus at a distance.
Bohr — Electrons in Shells
Niels Bohr proposed that electrons orbit the nucleus in specific, fixed energy levels (shells) rather than anywhere in the space around the nucleus. This explained the discrete emission spectra of elements. Each shell can hold a set number of electrons (the first holds up to 2, the second up to 8, and so on).
Chadwick — Discovery of the Neutron
James Chadwick discovered the neutron, confirming that the nucleus contains uncharged particles as well as protons. This explained why atoms are heavier than the number of protons alone would suggest, and why isotopes of the same element exist.
The plum pudding model predicted that all alpha particles would be deflected only slightly by the diffuse positive charge. Rutherford’s gold foil experiment showed that a small fraction bounced back sharply — something the plum pudding model could not explain. A new model was needed: the nuclear model, in which almost all the mass and all the positive charge are concentrated in a tiny nucleus at the centre of a largely empty atom. Rutherford’s result is a classic example of how a single experiment can overturn an established model in science.