How to Draw Orbital Box Diagrams Step by Step

An orbital box diagram (also called an arrows-in-boxes diagram or orbital diagram) shows exactly how the electrons in an atom are arranged: which orbitals are occupied, which electrons are paired, and how many are unpaired. It carries more information than a written configuration like 1s² 2s² 2p⁴, and it’s the clearest way to see why atoms behave as they do. This guide gives a reliable step-by-step method and plenty of worked examples.

What the diagram shows

• Each box (or line, or circle) represents one orbital.
• Each arrow represents one electron.
• An up arrow (↑) and a down arrow (↓) represent electrons with opposite spins.
• Boxes for the same subshell are drawn touching or grouped together; they have equal energy.

Subshell	Number of boxes
s	1
p	3
d	5
f	7

See shells vs subshells vs orbitals.

The three rules

1. Aufbau principle: fill orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p… See the Aufbau principle.
2. Pauli exclusion principle: a maximum of two electrons per box, with opposite spins (↑↓). See the Pauli exclusion principle.
3. Hund’s rule: within a subshell, put one electron in each box, all with the same spin, before pairing any. See Hund’s rule.

The method

1. Find the number of electrons. For a neutral atom, this equals the atomic number. For ions, adjust for the charge.
2. Write the electron configuration using the Aufbau order (e.g. 1s² 2s² 2p³).
3. Draw the boxes for each occupied subshell, in order of energy (1 box for s, 3 for p, 5 for d).
4. Fill each subshell:
   • For s: add ↑, then ↓.
   • For p, d, f: add one ↑ to each box first (Hund), then go back and add ↓ arrows.
5. Check: total arrows = number of electrons; no box has more than two arrows; no box has two arrows pointing the same way.

Worked examples

Hydrogen (1 electron): 1s¹

1s
↑

Helium (2 electrons): 1s²

1s
↑↓

Carbon (6 electrons): 1s² 2s² 2p²

1s	2s	2p	2p	2p
↑↓	↑↓	↑	↑

Two unpaired electrons in separate 2p orbitals, with parallel spins.

Nitrogen (7 electrons): 1s² 2s² 2p³

1s	2s	2p	2p	2p
↑↓	↑↓	↑	↑	↑

Three unpaired electrons: a half-filled p subshell.

Oxygen (8 electrons): 1s² 2s² 2p⁴

1s	2s	2p	2p	2p
↑↓	↑↓	↑↓	↑	↑

The fourth p electron pairs up in the first box; two unpaired electrons remain.

Neon (10 electrons): 1s² 2s² 2p⁶

1s	2s	2p	2p	2p
↑↓	↑↓	↑↓	↑↓	↑↓

All paired: a full shell, which explains neon’s lack of reactivity.

Phosphorus (15 electrons): 1s² 2s² 2p⁶ 3s² 3p³

1s	2s	2p	2p	2p	3s	3p	3p	3p
↑↓	↑↓	↑↓	↑↓	↑↓	↑↓	↑	↑	↑

Iron (26 electrons): [Ar] 3d⁶ 4s²

Using noble gas shorthand, you only need to draw the orbitals outside the argon core. See noble gas shorthand.

[Ar]	3d	3d	3d	3d	3d	4s
	↑↓	↑	↑	↑	↑	↑↓

Four unpaired electrons.

Diagrams for ions

1. Start from the neutral atom’s configuration.
2. Positive ions: remove electrons from the highest-n subshell first. For transition metals, remove 4s before 3d.
3. Negative ions: add electrons to the next available orbitals using the usual rules.

Example: Fe³⁺ (23 electrons)

Fe: [Ar] 3d⁶ 4s² → remove two 4s electrons and one 3d electron → Fe³⁺: [Ar] 3d⁵

[Ar]	3d	3d	3d	3d	3d	4s
	↑	↑	↑	↑	↑

Five unpaired electrons, a half-filled d subshell. This helps explain why Fe³⁺ is so stable in many compounds.

Example: O²⁻ (10 electrons)

O: 1s² 2s² 2p⁴ → add two electrons → O²⁻: 1s² 2s² 2p⁶, the same diagram as neon.

See electron configurations of ions.

Exceptions: chromium and copper

For a few elements, the actual ground state differs from the simple filling order:

Chromium: [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²)

[Ar]	3d	3d	3d	3d	3d	4s
	↑	↑	↑	↑	↑	↑

Six unpaired electrons.

Copper: [Ar] 3d¹⁰ 4s¹ (not 3d⁹ 4s²)

[Ar]	3d	3d	3d	3d	3d	4s
	↑↓	↑↓	↑↓	↑↓	↑↓	↑

See electron configuration exceptions.

What the diagram tells you

• Number of unpaired electrons, which predicts magnetic behaviour (paramagnetic if any are unpaired). See magnetic elements.
• Common bonding patterns: carbon’s four outer electrons (after promotion) or nitrogen’s three unpaired electrons relate to the number of bonds they typically form.
• Stability patterns, such as half-filled and filled subshells.

Orbital diagrams vs written configurations

A written configuration such as 1s² 2s² 2p⁴ tells you how many electrons are in each subshell. An orbital diagram adds how they’re arranged within each subshell: which orbitals hold pairs, which hold single electrons, and in which spin direction. For the s subshell there’s no difference, because it has only one orbital. For p, d and f subshells, the diagram carries extra information that the written form hides.

That extra information matters. Oxygen’s written configuration doesn’t tell you it has two unpaired electrons; its orbital diagram does. Iron’s configuration [Ar] 3d⁶ 4s² doesn’t show that four of the six d electrons are unpaired; the diagram makes it obvious. Whenever a question mentions unpaired electrons, magnetism or spin, draw the diagram.

Tips for neat diagrams

• Draw the boxes for a subshell touching each other, and leave a gap between subshells.
• Put subshells in order of energy from left to right (or from bottom to top if drawing an energy-level diagram).
• Use half-arrows or full arrows consistently, and draw the first electron in each box pointing up.
• For heavy elements, always use a noble gas core in brackets; nobody expects twenty boxes for the inner electrons.
• Label each group of boxes (1s, 2p, 3d) underneath so the marker can follow your reasoning.

A final check before handing in: count the arrows, count the unpaired electrons, and compare with the element’s position on the periodic table. A group 15 element should always show three unpaired p electrons.

Common mistakes

Mistake	Example	Correct version
Pairing before filling each box singly	C 2p: ↑↓, _, _	↑, ↑, _
Two same-spin arrows in one box	↑↑	↑↓
Too many arrows in a box	↑↓↑	maximum 2
Opposite spins for unpaired electrons	↑, ↓, _	↑, ↑, _
Wrong number of p or d boxes	two 2p boxes	three 2p boxes; five 3d boxes
Removing 3d before 4s for ions	Fe²⁺ as [Ar] 3d⁴ 4s²	[Ar] 3d⁶

Practice

Draw orbital box diagrams for:

1. Fluorine (Z = 9)
2. Sulfur (Z = 16)
3. Manganese (Z = 25)
4. Zn²⁺ (Z = 30)

Answers:

1. 1s ↑↓ | 2s ↑↓ | 2p ↑↓ ↑↓ ↑ → 1 unpaired
2. [Ne] 3s ↑↓ | 3p ↑↓ ↑ ↑ → 2 unpaired
3. [Ar] 3d ↑ ↑ ↑ ↑ ↑ | 4s ↑↓ → 5 unpaired
4. [Ar] 3d ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ (4s empty) → 0 unpaired

Check your answers with the electron configuration tool, which draws orbital boxes automatically.

Key takeaways

• Each box is an orbital, each arrow an electron; s, p, d and f have 1, 3, 5 and 7 boxes.
• Fill boxes in Aufbau order, with at most two opposite-spin arrows per box.
• Within a subshell, put one arrow in each box (same spin) before pairing.
• For transition metal ions, remove 4s electrons before 3d.
• Diagrams reveal unpaired electrons, magnetism and stability patterns at a glance.