The Periodic Table: Periodic Trends

Periodic Trends: Atomic Radius

Atomic radius is the distance from the centre of an atom’s nucleus to its outermost electron shell.
Across a period (left to right), the atomic radius generally decreases. This is because additional electrons are added to the same energy level while the nucleus increases in protons, causing stronger attraction between the nucleus and electrons.
Down a group (top to bottom), the atomic radius typically increases. This is due to the adding of new energy levels or electron shells, which are further from the nucleus.

Ionic radius is the measure of an atom’s ion in a crystal lattice.
Positive ions are smaller than the parent atoms as electron(s) are lost and repulsion between electrons decreases. Thus, cation radius is smaller than the atomic radius.
Negative ions are larger than the parent atoms as electron(s) are gained, resulting in increased repulsion between electrons. The anion radius is larger than the atomic radius.

Ionisation energy is the energy required to remove an electron from an atom.
Across a period from left to right, ionisation energy generally increases. This is due to the increased nuclear charge which strongly attracts the electrons in the same energy level.
Down a group from top to bottom, ionisation energy usually decreases. This is because outer electrons are further from the nucleus and are shielded by inner-level electrons, making them easier to remove.

Electronegativity is a measure of the ability of an atom to attract bonding electrons towards itself in a covalent bond.
Electronegativity typically increases across a period from left to right, due to the increased nuclear charge that more strongly attracts the bonding pair of electrons.
Electronegativity generally decreases down a group from top to bottom. This is due to the outer electrons being further from the nucleus and more shielded, reducing their attraction towards the bonding pair of electrons.

Electron affinity is the energy change when an atom gains an electron to form a negative ion.
Electron affinity generally increases from left to right across a period. This is due to the increased effective nuclear charge attracting additional electrons to the atom.
It generally decreases down a group. This is due to the outer electron shell getting further from the nucleus, reducing the attractive force on the new electron.