Which of the following shows the correct relationship between the first ionization energies of the given elements?

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BlackTom AI · Accepted Answer

Question restatement: Which of the following shows the correct relationship between the first ionization energies of the given elements? Option-by-option analysis: Option 1: Ne > F - This statement says the first ionization energy of Neon is greater than that of Fluorine. Across a period, ionization energy generally increases from left to right, and Neon (a noble gas with a complete octet) indeed has a higher first ionization energy than Fluorine. The actual values are roughly Ne ≈ 2080 kJ/mol and F ≈ 1680 kJ/mol, so Ne > F is correct based on periodic trends and empirical data. Option 2: Li > Be - Here we would be claiming Lithium has a higher first ionization energy than Beryllium. In the same period (period 2), Be is to the right of Li, and ionization energy increases across a period. Therefore Be should have a higher IE than Li. The correct relation is Li < Be, so this statement is false. Option 3: Ga > Ge - This compares Gallium (period 4, group 13) and Germanium (period 4, group 14). Within the same period, Ge, being to the right of Ga, generally has a higher ionization energy. In fact, Ga has IE ≈ 578 kJ/mol, while Ge has IE ≈ 762 kJ/mol, making Ga < Ge. The claim Ga > Ge is incorrect. Option 4: Xe > Ne - This option suggests Xenon has a higher first ionization energy than Neon. In fact, Neon has a higher IE because it is a noble gas with a complete valence shell and a relatively high effective nuclear charge compared to Xe, but Xe is larger and more easily ionized than Ne for the first electron due to shielding and energy level considerations. Empirically, Xe’s IE is about 1170 kJ/mol, whereas Ne’s is about 2080 kJ/mol, so Xe < Ne. Therefore this statement is false. Option 5: I am unsure and need to study this concept to prepare for the exam - This is not a statement about the actual ionization energies; it simply expresses uncertainty. While it’s okay to seek study help, it does not reflect a true comparison between the given elements. It’s not a valid evaluation of the relationship. Key takeaway for the correct option: The correct relationship is Ne > F because Neon’s first ionization energy is higher than Fluorine’s, in line with periodic trends and actual data. The other options fail to match known IE patterns across periods and within groups, or rely on incorrect assumptions about trends.

Which of the following shows the correct relationship between the first ionization energies of the given elements? 单项选择题

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