Why does the 4s orbital fill before the 3d orbital?

Due to the n+l rule, the 4s orbital (n=4, l=0, n+l=4) has a lower energy than the 3d orbital (n=3, l=2, n+l=5) for neutral atoms in their ground state. Therefore, electrons fill 4s first. However, once the 3d orbitals begin to fill (in transition metals), the 4s orbital's energy increases slightly, and electrons are typically lost from the 4s orbital before the 3d during ionization.

What do the up and down arrows in the orbitals represent?

The arrows represent electrons and their intrinsic spin, a quantum mechanical property. An up arrow denotes an electron with a spin quantum number of +1/2 (often called 'spin-up'), and a down arrow denotes -1/2 ('spin-down'). The Pauli exclusion principle requires that two electrons in the same orbital must have opposite spins, which is why they are always shown paired as an up-down pair.

Does this simulator show excited states or ions?

No, this model is specifically designed for building the ground-state electron configuration of neutral atoms. It follows the strict filling order of the Aufbau principle. To depict ions or excited states, electrons would need to be removed from or promoted to higher energy orbitals, which is beyond the scope of this simplified visualization.

Why do electrons fill orbitals singly before pairing up (Hund's rule)?

Electrons are negatively charged and repel each other. Occupying separate orbitals within the same subshell minimizes this repulsive electrostatic energy. Additionally, electrons have identical charges and prefer to have parallel spins (a consequence of quantum mechanics) when in different orbitals, which further stabilizes the arrangement. Pairing electrons in one orbital requires additional energy to overcome their mutual repulsion.

Electron Configuration

Fill orbitals visually with Aufbau principle animation.

Aufbau prediction

1s²2/2 e⁻

2s²2/2 e⁻

2p⁶6/6 e⁻

3s²2/2 e⁻

3p⁴4/6 e⁻

Reference (periodic dataset)

1s²2/2 e⁻

2s²2/2 e⁻

2p⁶6/6 e⁻

3s²2/2 e⁻

3p⁴4/6 e⁻

Compact notation

1s2 2s2 2p6 3s2 3p4

Sulfur or sulphur (see spelling differences) is a chemical element with symbol S and atomic number 16. It is an abundant, multivalent non-metal. Under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S8.

About this model

The arrangement of electrons within an atom, known as its electron configuration, is governed by a set of quantum mechanical rules. This simulator visualizes the core principles of this process: the Aufbau principle, Hund's rule, and the Pauli exclusion principle. The Aufbau principle (from the German for 'building up') dictates that electrons occupy the lowest energy atomic orbitals first. The simulator models this by filling orbitals in the order determined by the n+l rule, where 'n' is the principal quantum number and 'l' is the azimuthal quantum number. Orbitals with lower (n+l) values are filled first; if two orbitals have the same (n+l), the one with the lower 'n' is filled first. This yields the sequence 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc. Within a subshell (e.g., the three 2p orbitals), Hund's rule of maximum multiplicity is applied: electrons will fill degenerate orbitals singly and with parallel spins before pairing up. This is enforced to minimize electron-electron repulsion. The Pauli exclusion principle, a fundamental law of quantum mechanics, is represented by allowing a maximum of two electrons per orbital, and these two must have opposite spins. The simulator simplifies the complex quantum mechanical reality by representing orbitals as static boxes of fixed energy, ignoring the nuanced shapes of probability density clouds and the continuous nature of energy levels in multi-electron atoms. It also treats orbital energies as invariant, whereas in reality, the energy of an orbital can shift slightly as other orbitals are filled. By interacting with this model, students learn to predict ground-state electron configurations for the first 86 elements, understand the logic behind the periodic table's structure, and see the direct application of quantum rules to chemical behavior.

Who it's for: High school and introductory undergraduate chemistry students learning atomic structure and periodicity, as well as educators demonstrating quantum rules in a visual format.

Key terms

Aufbau Principle
Pauli Exclusion Principle
Hund's Rule
Atomic Orbital
Electron Spin
Quantum Number
Ground State
Electron Configuration

How it works

Aufbau filling (Klechkowski order) predicts how subshells fill with increasing Z. Many ground-state atoms follow it, but exceptions (Cr, Cu, Mo, Pd, …) lower the energy by promoting one s electron into d or f. The reference row uses curated configurations from the periodic dataset.

Key equations

Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → …

Frequently asked questions

Why does the 4s orbital fill before the 3d orbital?: Due to the n+l rule, the 4s orbital (n=4, l=0, n+l=4) has a lower energy than the 3d orbital (n=3, l=2, n+l=5) for neutral atoms in their ground state. Therefore, electrons fill 4s first. However, once the 3d orbitals begin to fill (in transition metals), the 4s orbital's energy increases slightly, and electrons are typically lost from the 4s orbital before the 3d during ionization.
What do the up and down arrows in the orbitals represent?: The arrows represent electrons and their intrinsic spin, a quantum mechanical property. An up arrow denotes an electron with a spin quantum number of +1/2 (often called 'spin-up'), and a down arrow denotes -1/2 ('spin-down'). The Pauli exclusion principle requires that two electrons in the same orbital must have opposite spins, which is why they are always shown paired as an up-down pair.
Does this simulator show excited states or ions?: No, this model is specifically designed for building the ground-state electron configuration of neutral atoms. It follows the strict filling order of the Aufbau principle. To depict ions or excited states, electrons would need to be removed from or promoted to higher energy orbitals, which is beyond the scope of this simplified visualization.
Why do electrons fill orbitals singly before pairing up (Hund's rule)?: Electrons are negatively charged and repel each other. Occupying separate orbitals within the same subshell minimizes this repulsive electrostatic energy. Additionally, electrons have identical charges and prefer to have parallel spins (a consequence of quantum mechanics) when in different orbitals, which further stabilizes the arrangement. Pairing electrons in one orbital requires additional energy to overcome their mutual repulsion.

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