Final answer:
The first ionization energy of sulfur is less than that of phosphorus due to the electron shielding effect and the repulsion between two electrons in the same p orbital, which makes the paired electron easier to remove.
Step-by-step explanation:
The reason why the first ionization energy of sulfur is less than that of phosphorus is mainly due to electron shielding effect (C). The ionization energy trend on the periodic table typically increases as we move from left to right across a period due to increasing nuclear charge. However, for sulfur, the added electron enters the 3p orbital, making it the first pair in that subshell. This causes electron-electron repulsion within the same orbital, leading to increased instability. As such, this paired electron is easier to remove, reducing the ionization energy for sulfur compared to phosphorus, which does not have paired p electrons.
Additionally, it is important to recognize that the shielding effect caused by the electrons in the 3s subshell does not significantly increase from phosphorus to sulfur to offset the additional proton in sulfur. Therefore, the choices A) Smaller atomic size of sulfur, B) Greater nuclear charge in sulfur, and D) Higher electronegativity of sulfur are not the primary reasons for the lower ionization energy in sulfur; instead, it is the shielding effect and electron repulsion that are crucial.