Ionisation Energy, Electron Affinity, and Electronegativity - SS2 Chemistry Lesson Note

Ionisation energy is the energy required to remove an electron from an atom or ion in the gaseous state. It is typically measured in units of kilojoules per mole (kJ/mol) or electron volts (eV). The first ionisation energy refers to the energy required to remove the outermost (valence) electron from a neutral atom, resulting in the formation of a positively charged ion (cation). Subsequent ionisation energies refer to the energy required to remove additional electrons from the ionised atom or ion. These energies are generally higher than the first ionisation energy due to the increased positive charge of the ion and the stronger attraction between the remaining electrons and the nucleus.

Ionisation energy generally follows periodic trends:

●     Across a period (from left to right), ionisation energy tends to increase. This is because the effective nuclear charge (positive charge experienced by the valence electrons) increases, resulting in a stronger attraction between the nucleus and the electrons.

●     Down a group (from top to bottom), ionisation energy generally decreases. This is primarily due to the increasing distance between the valence electrons and the nucleus, as the electron shells are added.

Electron Affinity:

Electron affinity is the energy change that occurs when an atom gains an electron to form a negatively charged ion (anion) in the gaseous state. It is typically expressed as a negative value since energy is released when an electron is added. Electron affinity can be influenced by factors such as atomic size, effective nuclear charge, and electron configuration. Electron affinity trends are not as consistent as ionisation energy trends, but some general patterns can be observed:

●     Across a period, electron affinity can increase or decrease. Elements on the right side of the periodic table tend to have higher electron affinities since they are closer to achieving a stable noble gas electron configuration.

●     Down a group, electron affinity generally decreases. This is due to the increasing atomic size and shielding effect, which reduces the effective nuclear charge and the attraction for additional electrons.

Electronegativity:

Electronegativity is a measure of the ability of an atom to attract electrons in a chemical bond. It is typically measured using various scales such as the Pauling scale or the Mulliken scale. Electronegativity values range from 0 to 4 or 0 to 4.0, depending on the scale used. Electronegativity trends generally follow the same patterns as ionisation energy:

●     Across a period, electronegativity tends to increase. This is because atoms on the right side of the periodic table have a stronger attraction for electrons, resulting in a higher electronegativity.

●     Down a group, electronegativity generally decreases. This is due to the increasing atomic size and the decreasing effective nuclear charge, which reduces the atom's ability to attract electrons.