Phosphorus has a higher first ionization energy than sulfur because sulfur’s fourth 3p electron is forced to share an orbital with another electron, creating repulsion that makes it easier to remove. Phosphorus has a first ionization energy of 10.49 eV, while sulfur’s is 10.36 eV. That difference is small, but it breaks the expected trend of steadily increasing ionization energy across Period 3, and the reason comes down to how electrons fill orbitals.
The General Trend Across Period 3
As you move from sodium to argon, ionization energy generally increases. Each element has one more proton in its nucleus than the last, which pulls the electrons in more tightly and makes them harder to remove. This steady buildup of nuclear charge is the dominant factor across the period.
But the trend isn’t perfectly smooth. There are two dips: one between magnesium and aluminum, and another between phosphorus and sulfur. Both dips happen for reasons related to how electrons are arranged in orbitals, not because the nuclear charge suddenly weakens.
How Phosphorus and Sulfur Fill Their Orbitals
Both elements have their outermost electrons in the 3p subshell, which contains three individual orbitals. Think of these as three separate slots that electrons fill one at a time before doubling up.
Phosphorus has the electron configuration 1s² 2s² 2p⁶ 3s² 3p³. Its three 3p electrons each occupy their own orbital, all spinning in the same direction. No orbital is shared. This is called a half-filled subshell, and it represents a particularly balanced, low-energy arrangement.
Sulfur has one more electron: 1s² 2s² 2p⁶ 3s² 3p⁴. Three of those four 3p electrons sit alone in their orbitals, but the fourth has to double up and share an orbital with one of the others. That pairing is the key to the whole anomaly.
Why Pairing Makes Removal Easier
Two electrons crammed into the same orbital are physically close to each other, and since both carry a negative charge, they repel one another. This repulsion raises the energy of both electrons, making one of them easier to knock loose. When you ionize sulfur, the electron you’re removing is that extra paired electron, and it’s already being pushed away by its orbital partner.
Phosphorus doesn’t have this problem. Every one of its 3p electrons sits alone, with no same-orbital repulsion to destabilize it. Even though sulfur has one more proton pulling its electrons inward, the added repulsion from pairing more than cancels out that extra nuclear attraction. The net result: it takes slightly less energy to remove an electron from sulfur than from phosphorus.
Effective Nuclear Charge Isn’t the Whole Story
You might expect sulfur’s higher nuclear charge (16 protons vs. 15 for phosphorus) to guarantee a higher ionization energy. Using Slater’s rules, a common method for estimating how much nuclear charge an outer electron actually “feels,” sulfur’s outermost electron does experience a slightly stronger pull than phosphorus’s. Same-shell electrons partially shield each other, while inner-shell electrons shield more effectively, but the net charge seen by sulfur’s outer electron is still marginally higher.
Under normal circumstances, that stronger pull would mean a higher ionization energy. But the electron-electron repulsion within that doubly occupied orbital is a competing effect, and in this case it wins. The repulsion penalty outweighs the benefit of one extra proton.
It’s worth noting that both elements have essentially the same atomic radius (about 180 pm by van der Waals measure), so distance from the nucleus isn’t a meaningful differentiator here. The explanation really does come down to what’s happening inside that one shared orbital.
The Same Pattern in Period 2
This isn’t unique to phosphorus and sulfur. The exact same anomaly shows up one row higher on the periodic table, between nitrogen and oxygen. Nitrogen has a half-filled 2p³ configuration with three unpaired electrons. Oxygen, at 2p⁴, forces one pair into the same orbital. The result: oxygen has a lower ionization energy than nitrogen, despite having a higher nuclear charge.
The pattern also repeats further down the periodic table with selenium, which sits below sulfur in Group 16. In every case, the element with a half-filled p subshell (Group 15) has a higher ionization energy than the element immediately to its right (Group 16). It’s one of the most consistent exceptions to the general left-to-right trend, and it appears in every period where p orbitals are being filled.
Half-Filled Stability in Context
Chemistry textbooks often describe half-filled subshells as “extra stable,” which is a useful shorthand but can be misleading. The half-filled configuration isn’t stable because of some special energetic bonus. Rather, it’s stable because it avoids the penalty of electron pairing. Every electron in phosphorus’s 3p subshell has its own orbital and spins in the same direction, minimizing repulsion. There’s no cost being paid for crowding.
At a deeper level, quantum mechanics explains this through something called exchange energy. Electrons with parallel spins (all pointing the same way) in different orbitals interact in a way that slightly lowers their total energy. A half-filled subshell maximizes the number of these favorable same-spin interactions. But for most purposes, the simpler explanation works perfectly well: pairing electrons in the same orbital costs energy because of repulsion, and phosphorus doesn’t have to pay that cost while sulfur does.