Atomic Structure and Periodicity - SS2 Chemistry Past Questions and Answers - page 6

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, the ionisation energy generally increases. This is because, as you move from left to right, 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, the ionisation energy generally decreases. This is because the atomic radius increases, and the shielding effect from inner electron shells reduces the effective nuclear charge. The increased distance and shielding decrease the attractive force on valence electrons, making it easier to remove an electron and resulting in lower ionisation energy.

The observed variations in ionisation energy are primarily determined by the balance between the attractive force of the protons in the nucleus and the repulsive force between electrons. The increasing effective nuclear charge and decreasing atomic radius across a period increase the attraction between the nucleus and valence electrons, making it more difficult to remove an electron. Conversely, the increasing atomic radius and shielding effect down a group weaken the attractive force, making it easier to remove an electron.

Electronegativity is the measure of an atom's ability to attract shared electrons in a chemical bond.

Across a period in the periodic table, electronegativity generally increases from left to right. This is due to the increasing effective nuclear charge and decreasing atomic radius, which result in a stronger attraction for shared electrons in a chemical bond.

Down a group, electronegativity generally decreases. This is because the atomic radius increases, leading to a decreased attraction for shared electrons. Additionally, the shielding effect from inner electron shells reduces the effective nuclear charge, further weakening the attractive force.

The observed variations in electronegativity are primarily influenced by the balance between the attractive force of the protons in the nucleus and the repulsive force between electrons. As you move across a period, the increasing effective nuclear charge enhances the attraction for shared electrons, resulting in higher electronegativity. Down a group, the increasing atomic radius and shielding effect weaken the attractive force, leading to lower electronegativity.

Electronegativity plays a vital role in determining the nature of chemical bonding and the polarity of molecules. Elements with higher electronegativity tend to attract electrons more strongly and form polar covalent or ionic bonds, while elements with lower electronegativity are more likely to form nonpolar covalent bonds.