Sulfur has a lower first ionization energy than phosphorus because sulfur has a paired electron in one of its 3p orbitals, and the repulsion between those two electrons makes one of them easier to remove. Phosphorus, by contrast, has exactly one electron in each of its three 3p orbitals, a half-filled arrangement that is unusually stable. The numbers reflect this: phosphorus requires 10.487 eV to remove its first electron, while sulfur requires only 10.360 eV.
This breaks the general rule that ionization energy increases as you move left to right across a period. Understanding why requires a closer look at how electrons are arranged inside these two atoms.
The General Trend and Why It Breaks
Across a row of the periodic table, each element has one more proton in its nucleus than the element before it. More protons means a stronger pull on the surrounding electrons, so it generally takes more energy to remove one. By that logic, sulfur (16 protons) should have a higher ionization energy than phosphorus (15 protons). But it doesn’t. The explanation lies entirely in how the outermost electrons are distributed among their orbitals.
How Phosphorus and Sulfur Fill Their Orbitals
Both phosphorus and sulfur have the same inner electron shells. The difference is in the 3p subshell, which contains three individual orbitals. Think of these orbitals as three separate rooms that electrons can occupy.
Phosphorus has three electrons in its 3p subshell. Following Hund’s rule, each electron occupies its own orbital, giving phosphorus one electron per room: ↑ | ↑ | ↑. No electron has to share space with another.
Sulfur has four electrons in its 3p subshell. The first three fill the three orbitals one each, but the fourth electron has no empty orbital left. It must pair up with an electron already in one of the orbitals: ↑↓ | ↑ | ↑. That paired orbital is where things get interesting.
Why Paired Electrons Are Easier to Remove
Two electrons sharing the same orbital are physically close to each other, and because they both carry a negative charge, they repel one another. This electron-electron repulsion raises the energy of both electrons in that orbital, making them less tightly bound to the atom. When you go to remove an electron from sulfur, you’re pulling away one of those repelling partners, which takes less effort than you’d expect from an atom with 16 protons.
In phosphorus, every 3p electron sits alone in its own orbital. There’s no partner pushing it away from the nucleus. Each electron feels the nuclear charge more directly, so each one is held more tightly.
The Stability of a Half-Filled Subshell
Phosphorus’s three unpaired 3p electrons represent a half-filled subshell, and this configuration is particularly stable for two reasons.
First, electrons in singly occupied orbitals are less effectively shielded from the nucleus. Shielding is what happens when inner electrons block the pull of the protons. When each orbital holds just one electron, the shielding effect is minimized, and each electron experiences a stronger effective nuclear charge.
Second, electrons with the same spin in different orbitals of the same subshell have a favorable quantum mechanical interaction called exchange energy. The more unpaired, same-spin electrons you have, the greater this stabilizing effect. Phosphorus, with three same-spin electrons across three orbitals, maximizes this benefit within the 3p subshell. Removing one of those electrons disrupts this arrangement, which costs extra energy.
So phosphorus gets a stability bonus that sulfur doesn’t. Sulfur’s fourth electron not only fails to gain the same exchange stabilization, it actively destabilizes the orbital it enters through repulsion.
Atomic Size Doesn’t Explain It
You might wonder whether the two atoms are different sizes, which could also affect ionization energy. They aren’t, at least not meaningfully. Both phosphorus and sulfur have a van der Waals radius of 180 picometers. The outermost electrons in both atoms sit at roughly the same distance from the nucleus. This confirms that the ionization energy difference isn’t about distance. It’s purely about the electron-electron repulsion in sulfur’s paired orbital and the half-filled stability of phosphorus.
Where This Exception Fits in the Periodic Table
This same type of exception appears elsewhere in the periodic table whenever an element’s electron configuration crosses from a half-filled to a more-than-half-filled subshell. Oxygen has a lower ionization energy than nitrogen for the exact same reason, just one row up: nitrogen has a half-filled 2p subshell (three unpaired electrons), while oxygen forces a fourth electron into a 2p orbital, creating the same pairing repulsion.
These dips don’t mean the overall trend is wrong. Ionization energy does generally increase across a period. But the transition from a half-filled to a partially paired subshell creates a predictable exception. Once you move past sulfur to chlorine and argon, the ionization energies resume their upward climb because the added nuclear charge more than compensates for any additional pairing.
The key takeaway is that two competing forces shape ionization energy: the pull of the nucleus and the repulsion between electrons sharing an orbital. In the case of phosphorus versus sulfur, electron repulsion wins.