Final answer:
Mg has a higher ionization energy than Sr, Cl has a higher ionization energy than S, and Br has a higher ionization energy than I. This is due to periodic trends where ionization energy increases across a period and decreases down a group, influenced by effective nuclear charge and electron shielding.
Step-by-step explanation:
Ionization energy refers to the amount of energy needed to remove an electron from an atom in its gaseous state. Among the given pairs, Mg has a higher ionization energy than Sr, Cl has a higher ionization energy than S, and Br has a higher ionization energy than I. This pattern is explained by the periodic table trends where ionization energy increases from left to right across a period due to increasing effective nuclear charge, and it generally decreases from top to bottom of a group as the outer electrons are further from the nucleus and more shielded by inner electrons.
For the pairs in question, Mg (3s2) is to the right of Sr (5s2) in the same group and Cl (3p5) is above S (3p4) in the same period. Similarly, Br (4p5) is above I (5p5) in the same group. The electron configuration plays a pivotal role in this trend. Mg's 3s electrons are closer to the nucleus and less shielded than Sr's 5s electrons. Cl's additional proton exerts a stronger pull on its electrons compared to S, making Cl's electrons harder to remove. Br is closer to the top of its group compared to I, therefore it experiences less shielding and a stronger attraction from the nucleus, leading to higher ionization energy