Oxygen (Group 16) has a slightly lower first ionization energy than nitrogen (Group 15), despite being further right in the same period. What is the best explanation?
AOxygen has more inner-shell electrons that shield its valence electrons more effectively than nitrogen's
BNitrogen's half-filled 2p subshell (three unpaired electrons) is unusually stable; oxygen's 2p has one paired electron whose electron-electron repulsion makes that electron easier to remove
COxygen's additional proton reduces effective nuclear charge by attracting inner electrons more strongly
DThe 2p electrons in oxygen are in a lower-energy orbital than those in nitrogen, requiring less energy to remove
Nitrogen has the electron configuration [He] 2s² 2p³ — three 2p electrons, one per orbital, all unpaired. No electron-electron repulsion occurs within the 2p subshell. Oxygen has [He] 2s² 2p⁴ — one 2p orbital must hold a pair. The repulsion between paired electrons in that orbital makes one of them easier to remove than any of nitrogen's unpaired 2p electrons, despite oxygen's higher nuclear charge. This is the Group 15–16 exception — one of two systematic dips in the otherwise rising IE trend across a period.
Question 2 Multiple Choice
Moving down Group 1 from Li to Cs, first ionization energy decreases even though nuclear charge increases substantially. Which factor is responsible?
AThe nucleus becomes less stable with more protons, weakening the nuclear force on outer electrons
BEach successive period adds a new electron shell, placing the valence electron farther from the nucleus and behind more inner-electron shielding — these effects outweigh the increase in nuclear charge
CEach Group 1 element has fewer total electrons than the previous one, reducing the nuclear attraction
DThe s orbital becomes more stable with increasing nuclear charge, paradoxically reducing ionization energy
Going down Group 1, nuclear charge increases, which would increase ionization energy in isolation. But each new period adds an entire new electron shell. The valence electron sits farther from the nucleus and behind a larger number of inner-shell electrons that shield it from the nuclear charge. Both the increased distance (the force weakens as 1/r²) and the increased shielding reduce the effective nuclear charge experienced by the outermost electron. The shielding and distance effects dominate, so ionization energy falls from Li to Cs.
Question 3 True / False
Boron (Group 13) has a lower first ionization energy than beryllium (Group 2), even though boron has one more proton.
TTrue
FFalse
Answer: True
Beryllium's outermost electron is in a 2s orbital. Boron's outermost electron is in a 2p orbital, which is higher in energy and has less penetration toward the nucleus than 2s. Even though boron has one more proton (higher Z), the electron being removed is in a less-stable, more-shielded orbital, making it easier to remove. This is the Group 2–13 exception — one of two systematic dips in the period trend.
Question 4 True / False
Ionization energy increases smoothly and without exception from left to right across nearly every period of the periodic table.
TTrue
FFalse
Answer: False
There are two notable dips in the otherwise rising trend across each period: (1) IE drops from Group 2 to Group 13 because the Group 13 element loses a p electron (higher energy than the s electron lost from Group 2); (2) IE drops from Group 15 to Group 16 because the Group 16 element's 2p subshell has a paired electron with added repulsion, making it easier to remove. These exceptions are not random — they reflect subshell structure and are reproducible across periods.
Question 5 Short Answer
Explain why sodium (Na) has a much lower first ionization energy than chlorine (Cl), even though both are in period 3.
Think about your answer, then reveal below.
Model answer: Both Na and Cl have their valence electrons in the n = 3 shell, but effective nuclear charge (Z_eff) differs dramatically. As protons are added from Na (Z = 11) to Cl (Z = 17) across period 3, electrons are added to the same 3s and 3p subshells. Electrons in the same shell shield each other poorly, so Z_eff rises steadily across the period. Na's single 3s electron experiences a Z_eff of roughly 2.5; Cl's 3p electrons experience a Z_eff of roughly 6. Cl's valence electrons feel far greater nuclear pull and require much more energy to remove.
This also explains why Na readily forms Na+ cations while Cl tends not to lose electrons. The periodic trend in ionization energy directly predicts electronegativity, reactivity, and ion formation tendencies. The same underlying logic — effective nuclear charge and electron shielding — governs all of these properties simultaneously.