Question
The first ionization energy of nitrogen (N) is higher than that of oxygen (O), even though oxygen comes after nitrogen in the periodic table. Explain this anomaly using electronic configurations.
Solution — Step by Step
First ionization energy (IE₁) is the energy required to remove one electron from a neutral gaseous atom:
General trend across Period 2: IE₁ increases from Li to Ne as you move left to right — more protons pull electrons in more tightly, so removing one takes more energy.
Expected order: Li < Be < B < C < N < O < F < Ne
Actual observed IE₁ values (kJ/mol): Li (520) < B (800) < Be (900) < C (1086) < O (1314) < N (1402) < F (1681) < Ne (2081)
Wait — N (1402) > O (1314)! The trend breaks between N and O. Why?
Nitrogen (Z = 7):
In the 2p subshell: three 2p orbitals, each with one electron (following Hund’s rule — all with same spin).
Oxygen (Z = 8):
In the 2p subshell: one orbital has two electrons paired.
Nitrogen’s 2p subshell is exactly half-filled (3 electrons in 3 orbitals). This configuration has special stability because:
- Symmetrical electron distribution — electrons are spread equally across all three 2p orbitals
- Exchange energy — electrons with the same spin in different orbitals can exchange positions; more same-spin electrons = more exchange energy = more stability
Nitrogen’s half-filled 2p subshell is more stable than expected. Removing an electron from this stable arrangement requires more energy than expected.
Oxygen has a paired electron in one 2p orbital. Paired electrons in the same orbital repel each other (electron-electron repulsion within the same small space). This inter-electronic repulsion makes one of those paired electrons easier to remove — it is already “destabilised” by repulsion.
When oxygen loses one electron, the pairing is broken:
This actually gives the stable half-filled configuration — so the system is moving toward stability upon ionisation.
| Species | 2p configuration | Reason for relative IE |
|---|---|---|
| N | (half-filled) | Extra stable → harder to remove electron → higher IE |
| O | (paired electron) | Pairing repulsion → easier to remove paired electron → lower IE |
N has higher first IE than O because nitrogen’s half-filled 2p subshell is exceptionally stable (due to exchange energy and symmetry), while oxygen’s paired electron in 2p is destabilised by electron-electron repulsion.
Why This Works
This is a recurring pattern across the periodic table. Any time a subshell is exactly half-filled or completely filled, that configuration has extra stability. The same anomaly appears between:
- Be and B — Be has a full 2s subshell (); extra stability makes IE(Be) > IE(B), even though B comes after Be
- Mg and Al — same reason, full 3s vs entering 3p
All these anomalies follow the same rule: half-filled and completely-filled subshells are extra stable → harder to ionise → higher IE than the next element.
Alternative Method — Comparing Shielding
Another way to think about it: when oxygen adds its 8th electron into an already-occupied 2p orbital, that electron experiences increased electron-electron repulsion (shielding from the other electron in the same orbital). This effectively reduces the nuclear attraction felt by that electron, making it easier to remove.
Nitrogen’s electrons are all in separate orbitals — no such extra shielding/repulsion within the 2p subshell.
This anomaly (N > O in IE) is one of the most frequently tested concepts in JEE Main and NEET. The explanation is always the same: half-filled subshell stability of nitrogen. Memorise: “N is half-filled → extra stable → higher IE than O despite O having more protons.”
Common Mistake
Students sometimes explain this anomaly by saying “oxygen has more protons, so it should have higher IE.” This is the general trend expectation — but that’s exactly why this IS an anomaly. The electronic structure effect (half-filled stability, pairing repulsion) outweighs the nuclear charge effect here. Always mention BOTH the reason N is extra stable AND why O is easier to ionise (paired electron repulsion) for full marks in CBSE/JEE answers.