Key Concepts: History of the Atom

Early concepts of the atom were proposed by Dalton in the 1800s, who suggested atoms were indivisible, solid spheres.
The ‘plum pudding’ model was suggested by J.J. Thompson in the late 1890s. He proposed that atoms were spheres of a positive ‘pudding’ with negative ‘plums’ (electrons) embedded throughout.
An experiment by Ernest Rutherford in 1909, involving firing positively charged alpha particles at a thin sheet of gold foil, led to the discovery of the nuclear model. Most particles went straight through, but some deflected or bounced back, implying a tiny, dense, positive nucleus at the centre of the atom.
Niels Bohr further refined the nuclear model in 1913. He suggested electrons exist in fixed orbits, or energy levels, around the nucleus. This is most similar to the concept of atom held today.
The modern model of the atom, known as the quantum mechanical model, explains phenomena that earlier models couldn’t, such as the emission spectrum of each element. It suggests that the exact position of an electron around a nucleus can never be precisely pinned down, instead we predict the area it’s most likely to be in - an ‘electron cloud’.
James Chadwick in 1932 discovered the neutron, a particle with no charge that exists in the nucleus of the atom. This discovery helped explain isotopes - atoms of the same element with different numbers of neutrons.
Atoms for each element can be represented using atomic and mass numbers. The atomic number, or proton number, is the number of protons in an atom of an element. The mass number is the sum of protons and neutrons in the atomic nucleus.
Isotopes, while having different numbers of neutrons, do not differ in chemical reactions due to the fact these interactions involve only electrons.
The position of an element on the periodic table provides significant information about its properties. Elements are organised by increasing atomic number. Each row (period) represents a new energy level for electrons, while each column (group) contains elements with the same number of electrons in their outer shell.
The periodic table reveals patterns of reactivity. For example, elements in Group 1 (alkali metals) are highly reactive, while elements in Group 18 (noble gases) are unreactive due to their full outer electron shells.
Knowledge of atomic structure helps in understanding chemical bonds - atoms gain, lose, or share electrons in their outer shell to achieve a stable electron configuration, leading to ionic or covalent bonds.
Ionisation energies (the energy required to remove an electron from an atom) increase across a period as the atomic radius decreases, and decrease down a group as the atomic radius increases. This is due to the effect of nuclear charge and electron shielding.
Understanding atomic structure and the periodic table are crucial for recognising and predicting trends in chemical reactions and behaviours of elements and compounds.