The orbital diagram for oxygen shows 8 electrons distributed across three subshells: 1s, 2s, and 2p. Its electron configuration is 1s²2s²2p⁴, meaning two electrons fill the first energy level and six electrons occupy the second. The key detail most students need to understand is what happens in the 2p subshell, where four electrons spread across three available orbitals following specific rules.
How the 8 Electrons Fill Each Orbital
An orbital diagram uses boxes (or horizontal lines) to represent individual orbitals, with arrows inside them representing electrons. An up arrow and a down arrow in the same box mean two electrons are paired with opposite spins. Here is how oxygen’s 8 electrons fill in order from lowest to highest energy:
1s orbital: One box, two arrows (↑↓). The first two electrons pair up here, completely filling the lowest energy level.
2s orbital: One box, two arrows (↑↓). The next two electrons pair up in the 2s orbital, which is the next step up in energy.
2p orbitals: Three boxes side by side, representing the three 2p orbitals (sometimes labeled 2px, 2py, and 2pz). The remaining four electrons go here, and this is where the diagram gets interesting. The first three electrons each go into a separate box with their spins pointing the same direction (↑)(↑)(↑). Only after each box has one electron does the fourth electron double up in the first box, pairing with opposite spin: (↑↓)(↑)(↑).
The complete orbital diagram looks like this:
1s: [↑↓] 2s: [↑↓] 2p: [↑↓] [↑] [↑]
Why the 2p Electrons Spread Out First
Three rules govern how electrons fill orbitals, and all three are visible in oxygen’s diagram.
The Aufbau principle says electrons fill lower-energy orbitals before higher-energy ones. That’s why the 1s fills before the 2s, and the 2s fills before the 2p. For oxygen, the energy order is straightforward: 1s → 2s → 2p.
The Pauli exclusion principle says each orbital can hold a maximum of two electrons, and those two must have opposite spins. That’s why every filled box in the diagram has one up arrow and one down arrow, never two pointing the same way.
Hund’s rule is the one that matters most for the 2p subshell. It says that when you have multiple orbitals at the same energy level (like the three 2p orbitals), electrons will each occupy their own orbital before any pairing occurs. Think of it like passengers on a bus: people tend to sit in empty seats before doubling up. The first three of oxygen’s four 2p electrons each claim their own orbital. Only the fourth is forced to pair up.
Oxygen Has Two Unpaired Electrons
Looking at the completed diagram, you can count two unpaired electrons in the 2p subshell. Those two boxes with single up arrows are what make atomic oxygen paramagnetic, meaning it is attracted to a magnetic field. If all electrons were paired, the atom would be diamagnetic and slightly repelled by magnets instead.
The number of unpaired electrons also explains why oxygen is so chemically reactive. Those two unpaired electrons are essentially “looking” for partners, which is why oxygen readily forms two covalent bonds with other atoms. In water (H₂O), for example, two hydrogen atoms each share an electron with oxygen, pairing up both of those lone electrons.
Valence Electrons and the Outer Shell
Oxygen’s first shell (1s²) holds 2 electrons, and its second shell (2s²2p⁴) holds 6. Those 6 electrons in the second shell are the valence electrons, the ones involved in chemical bonding. When you see oxygen’s orbital diagram, the 1s electrons are essentially buried in the core of the atom and don’t participate in reactions. Everything interesting happens in the 2s and 2p orbitals.
Because oxygen has 6 valence electrons and needs 8 to complete its outer shell (the octet rule), it tends to gain or share 2 electrons in chemical reactions. This is directly visible in the orbital diagram: two of the three 2p orbitals are only half-filled, giving oxygen the capacity to form two bonds.
Atomic Oxygen vs. Molecular Oxygen (O₂)
The orbital diagram described above is for a single oxygen atom. Molecular oxygen, the O₂ you breathe, has its own molecular orbital diagram that looks quite different. When two oxygen atoms bond together, their atomic orbitals combine to form new molecular orbitals that belong to the molecule as a whole.
What’s notable about O₂ is that its molecular orbital diagram also produces unpaired electrons, two of them sitting in separate antibonding orbitals. This is why liquid oxygen is paramagnetic and can be suspended between the poles of a strong magnet. Early chemists found this surprising because simple bonding models predicted all electrons in O₂ should be paired. The molecular orbital diagram was one of the first tools that correctly explained this magnetic behavior.
For most chemistry courses, when someone asks for “the orbital diagram for oxygen,” they mean the atomic version: the 1s²2s²2p⁴ box diagram for a single oxygen atom with its two unpaired electrons in the 2p subshell.