Aufbau, Pauli and Hund's Rule: How Electrons Fill Orbitals

Each rule answers a different question. Aufbau answers WHICH orbital fills first: always the lowest energy one, found using the (n+l) rule. Pauli answers HOW MANY electrons an orbital can hold: at most 2, and they must spin opposite ways. Hund answers HOW to fill orbitals of equal energy: place one electron in each orbital first (all same spin), then start pairing. Think of it as order, then limit, then arrangement.

For any orbital, add its principal quantum number n and azimuthal quantum number l. The orbital with the smaller (n+l) value has lower energy and fills first. If two orbitals have the same (n+l) value, the one with the smaller n fills first. Example: 4s has n+l = 4+0 = 4, and 3d has n+l = 3+2 = 5. Since 4 is smaller, 4s fills before 3d. This is why 4s comes before 3d even though 3 is less than 4.

4s fills first. Using the (n+l) rule: 4s gives 4+0 = 4 and 3d gives 3+2 = 5. Lower (n+l) means lower energy, so 4s (value 4) is filled before 3d (value 5). This is a very common NEET trap because 3 looks smaller than 4, but the rule is about n+l, not just n.

No. This is the Pauli Exclusion Principle. No two electrons in one atom can have all four quantum numbers the same. Two electrons in the same orbital already share n, l and m (magnetic). So they MUST differ in the fourth one, the spin quantum number: one is +1/2 and the other is -1/2. That is why an orbital holds a maximum of 2 electrons, and they always spin opposite.

This is Hund's rule of maximum multiplicity. When orbitals have the same energy (like the three 2p orbitals), electrons spread out one per orbital, all with parallel (same) spin, before any pairing. They do this because electrons repel each other. Staying in separate orbitals keeps them farther apart, which lowers energy and gives extra stability from something called exchange energy. So nitrogen is 2p to the power one in each of px, py, pz, not two electrons in one p orbital.

Multiplicity is linked to the number of unpaired electrons with parallel spin. 'Maximum multiplicity' means the correct ground-state arrangement is the one with the greatest number of unpaired electrons of the same spin. Filling each equal-energy orbital singly first gives the most unpaired electrons, so that arrangement is chosen. This matters in NEET for counting unpaired electrons and calculating magnetic moment.

The rules give the ground-state configuration, but a few elements look like exceptions because half-filled and fully-filled subshells are extra stable. Chromium is 3d to the power 5, 4s to the power 1 (not 3d to the power 4, 4s to the power 2) and copper is 3d to the power 10, 4s to the power 1 (not 3d to the power 9, 4s to the power 2). One 4s electron shifts to 3d to make it exactly half-filled or fully-filled, which lowers energy. This is a stability effect, not a true violation of Aufbau.

First use Aufbau to pick the next lowest-energy orbital (n+l rule). Then apply Pauli, so each orbital takes at most 2 electrons with opposite spins. Within a set of equal-energy orbitals, apply Hund, filling them singly with parallel spin before pairing. In practice all three work together for each subshell you fill.

4s fills before 3d because it is lower in energy in the neutral atom (n+l rule). But once 3d is occupied, 3d drops below 4s in energy. So when forming cations of transition metals, electrons are removed from 4s first. This is a common NEET point for d-block ions.

No, an s subshell has only one orbital, so there is nothing to spread across. Hund's rule matters only where a subshell has more than one equal-energy orbital: p (3 orbitals), d (5 orbitals) and f (7 orbitals).

One orbital holds 2 electrons (Pauli). A subshell holds 2(2l+1): s holds 2, p holds 6, d holds 10, f holds 14. A shell holds 2n² electrons, where n is the shell number.

Exchange energy is a stabilising energy that grows when more electrons of the same spin are present in a subshell and can swap positions. Filling orbitals singly with parallel spins (Hund's way) maximises exchange energy, which lowers the atom's energy and explains the extra stability of half-filled and fully-filled subshells.