The ionization energy is the energy required to remove an electron from a neutral atom, resulting in the formation of a positively charged ion. In the case of sulfur (S) and phosphorus (P), sulfur has a higher ionization energy than phosphorus due to a few factors:
Effective nuclear charge: The ionization energy depends on the attractive force between the positively charged nucleus and the negatively charged electron. Sulfur has a greater effective nuclear charge than phosphorus because it has more protons in its nucleus. The additional protons in sulfur's nucleus exert a stronger pull on the electrons, making it more difficult to remove an electron and therefore increasing its ionization energy.
Electron shielding: Electrons in an atom are distributed in different energy levels or shells. The inner electrons shield the outer electrons from the full force of the positive nucleus. In sulfur, there are more inner electron shells compared to phosphorus, which leads to increased electron-electron repulsion and less effective shielding of the outermost electrons. As a result, the outermost electron in sulfur experiences a stronger attraction to the nucleus, requiring more energy to remove it and resulting in a higher ionization energy.
Electron configuration: Sulfur and phosphorus belong to different periods on the periodic table, so they have different electron configurations. Sulfur has a configuration of 1s² 2s² 2p⁶ 3s² 3p⁴, while phosphorus has a configuration of 1s² 2s² 2p⁶ 3s² 3p³. In sulfur, the outermost electron is in a p-orbital, which is closer to the nucleus and more tightly held compared to the s-orbital in phosphorus. The p-electrons experience greater effective nuclear charge and less shielding, resulting in a higher ionization energy compared to phosphorus.
Overall, the combination of a greater effective nuclear charge, reduced electron shielding, and specific electron configurations contributes to sulfur having a higher ionization energy than phosphorus.