Atomic Structure and Periodicity - SS2 Chemistry Past Questions and Answers - page 4

Ionisation energy refers to the energy required to remove an electron from an atom or ion in the gaseous state.

Across a period in the periodic table, ionisation energy generally increases. This is because, as you move from left to right across a period, the atomic radius decreases, and the effective nuclear charge (the net positive charge experienced by valence electrons) increases. The increased attraction between the positively charged nucleus and the electrons requires more energy to remove an electron, leading to higher ionisation energy.

Down a group in the periodic table, ionisation energy generally decreases. This is due to the increasing atomic radius and the shielding effect from inner electron shells. The larger atomic radius and shielding reduce the effective nuclear charge, making it easier to remove an electron and resulting in lower ionisation energy.

Electron affinity refers to the energy change that occurs when an atom or ion in the gaseous state gains an electron to form a negatively charged ion.

Across a period in the periodic table, electron affinity does not have a consistent trend. However, there is a general increase in electron affinity as you move from left to right across a period. This is because, as the atomic radius decreases and effective nuclear charge increases, the attraction between the nucleus and incoming electrons strengthens. Therefore, atoms across a period have a higher tendency to gain electrons and exhibit higher electron affinity.

Down a group in the periodic table, electron affinity generally decreases. This is mainly due to the increasing atomic radius and the shielding effect. The larger atomic radius and shielding reduce the effective nuclear charge, weakening the attraction between the nucleus and incoming electrons. As a result, atoms down a group have a lower tendency to gain electrons and exhibit lower electron affinity.

Electronegativity refers to an atom's ability to attract electrons in a chemical bond.

Across a period in the periodic table, electronegativity generally increases. This is because, as you move from left to right, the atomic radius decreases, and the effective nuclear charge increases. The decreased atomic radius brings valence electrons closer to the nucleus, resulting in a stronger attraction and higher electronegativity. Atoms across a period have a greater ability to attract electrons and form polar covalent or ionic bonds.

Down a group in the periodic table, electronegativity generally decreases. This is due to the increasing atomic radius and the shielding effect from inner electron shells. The larger atomic radius and shielding reduce the effective nuclear charge, weakening the attractive force on electrons. Consequently, atoms down a group have a lower ability to attract electrons and form polar covalent or ionic bonds.

Ionisation energy and electron affinity are both related to the behaviour of electrons in atoms, but they represent different processes.

Ionisation energy refers to the energy required to remove an electron from an atom or ion in the gaseous state. It measures the energy needed to overcome the attractive forces between the electron and the nucleus. Higher ionisation energy indicates a stronger hold on electrons and a higher tendency to lose electrons.

Electron affinity, on the other hand, refers to the energy change that occurs when an atom or ion in the gaseous state gains an electron to form a negatively charged ion. It measures the energy released or absorbed when an electron is added to an atom. Higher electron affinity indicates a stronger attraction for electrons and a higher tendency to gain electrons.