Nitrogen (N) has a larger ionization energy compared to phosphorus (P), sulfur (S), and arsenic (As) due to several factors, including atomic structure and effective nuclear charge.
Firstly, ionization energy generally increases across a period from left to right in the periodic table. As you move from left to right, the atomic number and the number of protons in the nucleus increase. This results in a stronger electrostatic attraction between the positively charged nucleus and the negatively charged electrons, making it more difficult to remove an electron. Hence, nitrogen, being to the left of phosphorus, sulfur, and arsenic, experiences a higher ionization energy.
Moreover, nitrogen has a half-filled p orbital in its valence shell. This electron configuration provides added stability due to the exchange energy associated with electrons occupying degenerate (equal energy) orbitals. The half-filled p orbital makes it energetically unfavorable to remove an electron, leading to a higher ionization energy compared to the elements that follow in the same group.
Additionally, effective nuclear charge plays a role in determining ionization energy. Effective nuclear charge refers to the net positive charge experienced by the valence electrons and depends on the shielding effect of inner electrons. While nitrogen and the other elements in question have the same number of valence electrons (five), nitrogen has fewer inner electrons to shield the valence electrons from the attraction of the nucleus. As a result, the valence electrons in nitrogen experience a stronger effective nuclear charge, making it more difficult to remove an electron and resulting in a higher ionization energy.
In summary, nitrogen (N) has a larger ionization energy compared to phosphorus (P), sulfur (S), and arsenic (As) due to factors such as atomic structure, the stability of a half-filled p orbital, and the higher effective nuclear charge experienced by its valence electrons.