Provide a brief explanation for each of the following: (a) Cl^- is larger than Ar. (b) P^3- is larger than S^2-. (c) K⁺ is larger than Na⁺. (d) F^- is larger than F.

Verified step by step guidance

1. Consider the position of K and Na in the periodic table. K (potassium) is located in period 4, while Na (sodium) is in period 3.
2. Understand that as you move down a group in the periodic table, the number of electron shells increases. This means K has more electron shells than Na.
3. Recognize that both K+ and Na+ ions have lost one electron compared to their neutral atoms, resulting in the same electron configuration as the noble gas preceding them.
4. Despite having the same electron configuration, K+ has more electron shells than Na+, making it larger in size.
5. Conclude that the increase in the number of electron shells as you move down a group results in a larger ionic radius for K+ compared to Na+.

Key Concepts

Ionic Radius

Ionic radius refers to the size of an ion in a crystal lattice. Cations, like K⁺ and Na⁺, are formed when atoms lose electrons, resulting in a decrease in size due to reduced electron-electron repulsion and increased effective nuclear charge. The ionic radius can vary based on the ion's charge and the number of electron shells.

Periodic Trends

Periodic trends are patterns observed in the properties of elements across the periodic table. As you move down a group, the ionic radius increases due to the addition of electron shells, which outweighs the increase in nuclear charge. This explains why K⁺, located below Na⁺ in Group 1, is larger.

Effective Nuclear Charge

Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in a multi-electron atom. It accounts for the shielding effect of inner electrons. In the case of K⁺ and Na⁺, although both are cations, K⁺ has more electron shells, leading to a larger ionic radius despite a similar Z_eff.