Use electron configurations to explain the following observations: (a) The first ionization energy of phosphorus is greater than that of sulfur. (b) The electron affinity of nitrogen is lower (less negative) than those of both carbon and oxygen. (c) The second ionization energy of oxygen is greater than the first ionization energy of fluorine. (d) The third ionization energy of manganese is greater than those of both chromium and iron.

Verified step by step guidance

1. Step 1: Understand the concept of electron configuration and how it relates to ionization energy and electron affinity. Ionization energy is the energy required to remove an electron from an atom, while electron affinity is the energy change when an electron is added to an atom.
2. Step 2: For part (a), compare the electron configurations of phosphorus (P) and sulfur (S). Phosphorus has the configuration [Ne] 3s² 3p³, and sulfur has [Ne] 3s² 3p⁴. The half-filled p subshell in phosphorus is more stable, making it harder to remove an electron compared to sulfur.
3. Step 3: For part (b), examine the electron configurations of nitrogen (N), carbon (C), and oxygen (O). Nitrogen has [He] 2s² 2p³, carbon has [He] 2s² 2p², and oxygen has [He] 2s² 2p⁴. The half-filled p subshell in nitrogen is relatively stable, resulting in a lower electron affinity compared to carbon and oxygen.
4. Step 4: For part (c), consider the electron configurations of oxygen (O) and fluorine (F). Oxygen's configuration is [He] 2s² 2p⁴, and fluorine's is [He] 2s² 2p⁵. The second ionization energy of oxygen involves removing an electron from a positively charged ion, which requires more energy than removing an electron from a neutral fluorine atom.
5. Step 5: For part (d), analyze the electron configurations of manganese (Mn), chromium (Cr), and iron (Fe). Manganese has [Ar] 3d⁵ 4s², chromium has [Ar] 3d⁵ 4s¹, and iron has [Ar] 3d⁶ 4s². The third ionization energy of manganese involves removing an electron from a half-filled d subshell, which is more stable and requires more energy compared to chromium and iron.

Key Concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It generally increases across a period due to increasing nuclear charge, which holds electrons more tightly. However, it can vary due to electron-electron repulsions and the stability of electron configurations, such as full or half-full subshells.

Electron Affinity

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom in the gas phase. A more negative electron affinity indicates a greater tendency to gain an electron. Trends in electron affinity can be influenced by atomic size and the effective nuclear charge, affecting how tightly an atom can hold onto additional electrons.

Comparative Ionization Energies

Comparative ionization energies involve analyzing the energy required to remove successive electrons from an atom. The second ionization energy is often higher than the first due to the increased positive charge of the ion, which attracts remaining electrons more strongly. Understanding these trends helps explain why certain elements exhibit higher ionization energies than others, based on their electron configurations.