How well can you apply electron configuration rules? Work through 20 AP Chemistry-style practice questions covering orbitals, ions, valence electrons, and electron filling.
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The Pre-Quiz Review
Electron configuration questions become much easier when you know the filling order, the rules governing orbital occupation, and how configurations change when atoms form ions.
1. Aufbau Principle
The Aufbau principle states that electrons occupy available orbitals in order of increasing energy. To write an electron configuration, fill the lower-energy orbitals before moving to higher-energy orbitals.
Filling Order
The orbital sequence to remember
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
The order is based on relative orbital energies rather than simply increasing principal energy level. Notice, for example, that 4s is filled before 3d in the ground-state configurations of neutral atoms.
Use the diagonal rule
A common way to memorize the filling order is to write the orbitals in rows by principal energy level and follow diagonal arrows from the upper rows toward the lower rows. This produces the sequence above.
Orbital Capacities
An s subshell contains 1 orbital and holds up to 2 electrons. A p subshell contains 3 orbitals and holds up to 6 electrons. A d subshell contains 5 orbitals and holds up to 10 electrons, while an f subshell contains 7 orbitals and holds up to 14 electrons.
Capacity Pattern
The maximum capacities are s², p⁶, d¹⁰, and f¹⁴. These values are useful when checking whether an electron configuration is possible.
Example:
Write the configuration for phosphorus
Phosphorus has atomic number 15, so a neutral phosphorus atom contains 15 electrons.
P: 1s² 2s² 2p⁶ 3s² 3p³
The superscripts add to 15: 2 + 2 + 6 + 2 + 3 = 15.
The superscripts in a neutral atom’s electron configuration should add up to its atomic number.
2. Pauli Exclusion Principle
The Pauli exclusion principle states that an orbital can contain a maximum of two electrons, and those two electrons must have opposite spins.
One Orbital
↑↓ – allowed
Two electrons can occupy the same orbital when their spins are opposite.
Not Allowed
↑↑ – not allowed
Two electrons in the same orbital cannot have identical spin quantum states.
Why can an s subshell hold only two electrons?
An s subshell contains only one orbital. Since each orbital can hold a maximum of two electrons, the s subshell has a maximum capacity of two electrons.
s = 1 orbital × 2 electrons = 2 electrons
3. Hund’s Rule
Hund’s rule states that electrons occupy degenerate orbitals individually before pairing. In other words, electrons spread out across orbitals of equal energy before two electrons occupy the same orbital.
Here are two examples of Hund’s Rule of Maximum Multiplicity:
Carbon (C): In its 1s² 2s² 2p² configuration, the two 2p electrons occupy separate orbitals singly with parallel spins rather than pairing up in a single orbital.
Nitrogen (N): In its 1s² 2s² 2p³ configuration, all three 2p electrons individually occupy separate orbitals with parallel spins to maximize stability before any pairing begins.
Hund’s rule is especially important when determining the number of unpaired electrons and whether a species is paramagnetic.
4. Electron Configuration Notation
Electron configurations can be represented in several related ways. AP Chemistry questions may require you to recognize or move between these representations.
Full Configuration
Shows every occupied subshell.
Na: 1s² 2s² 2p⁶ 3s¹
Noble-Gas Notation
Replaces the inner electrons with the preceding noble gas.
Na: [Ne] 3s¹
Example
Iron
Iron has atomic number 26.
Fe: [Ar] 4s² 3d⁶
The noble-gas core accounts for 18 electrons, while the remaining 8 electrons occupy the 4s and 3d subshells.
5. Valence Electrons and Core Electrons
Valence electrons are the electrons most directly involved in chemical behavior. Core electrons are the inner electrons that generally do not participate directly in bonding.
Chlorine
Cl: [Ne] 3s² 3p⁵
Chlorine has seven electrons in its outer principal energy level, n = 3. These are its valence electrons.
For many main-group atoms, the number of valence electrons can be connected directly to the element’s periodic-table group.
6. Electron Configurations of Ions
Forming an ion changes the number of electrons, not the number of protons. Cations lose electrons, while anions gain electrons.
- Start with the neutral atom’s electron configuration.
- Determine whether electrons must be added or removed.
- Add electrons to form an anion or remove electrons to form a cation.
- Check that the final number of electrons matches the charge.
Sodium ion
Na: [Ne] 3s¹
Na⁺: [Ne]
Sodium loses one electron, leaving it with the same number of electrons as neon.
Chloride ion
Cl: [Ne] 3s² 3p⁵
Cl⁻: [Ne] 3s² 3p⁶
Chlorine gains one electron to complete its 3p subshell.
Transition-metal cations
When transition metals form cations, electrons are removed from the 4s orbital before the 3d orbitals, even though 4s is filled before 3d in the neutral atom’s ground-state filling order.
Iron(II) and iron(III)
Fe: [Ar] 4s² 3d⁶
Fe²⁺: [Ar] 3d⁶
Fe³⁺: [Ar] 3d⁵
The two 4s electrons are removed first. The third electron removed for Fe³⁺ comes from the 3d subshell.
7. Common Electron Configuration Exceptions
The Aufbau filling order is an extremely useful guide, but a few atoms have experimentally observed ground-state configurations that differ from the simplest predicted pattern.
| Element | Common predicted pattern | Observed configuration |
|---|---|---|
| Cr | [Ar] 4s² 3d⁴ | [Ar] 4s¹ 3d⁵ |
| Cu | [Ar] 4s² 3d⁹ | [Ar] 4s¹ 3d¹⁰ |
These configurations are often taught as exceptions because the observed arrangements provide additional stability associated with the d-subshell electron distribution.
8. Unpaired Electrons and Paramagnetism
A species is paramagnetic when it contains one or more unpaired electrons. A species with all electrons paired is diamagnetic.
Paramagnetic
At least one electron remains unpaired.
Diamagnetic
All electrons are paired.
Aluminum
Al: [Ne] 3s² 3p¹
The single 3p electron is unpaired, so a ground-state aluminum atom is paramagnetic.
9. AP Chemistry Quick Check
- Aufbau: fill orbitals in increasing energy order.
- Pauli: each orbital holds a maximum of two electrons with opposite spins.
- Hund: fill degenerate orbitals singly before pairing.
- s, p, d, f capacities: 2, 6, 10, and 14 electrons.
- Ions: cations lose electrons, and anions gain electrons.
- Transition-metal cations: remove 4s electrons before 3d electrons.
- Paramagnetic: at least one unpaired electron.
Quiz Tip
When solving an electron-configuration problem, determine the total number of electrons first, fill the orbitals using the appropriate rules, and then inspect the resulting configuration for valence electrons, unpaired electrons, ion formation, or periodic trends.
Watch Out: Anomalies in the d-Block
Transition-metal electron configurations are a common AP Chemistry trap because a few elements do not follow the simplest Aufbau prediction. Chromium (Cr) is the classic example.
Based on straightforward filling, you might expect Cr to be [Ar] 4s² 3d⁴. Instead, its ground-state configuration is [Ar] 4s¹ 3d⁵. This arrangement gives the 3d subshell five unpaired electrons, a particularly stable half-filled subshell.
Copper (Cu) shows a similar pattern: instead of [Ar] 4s² 3d⁹, its actual configuration is [Ar] 4s¹ 3d¹⁰, giving a completely filled 3d subshell.
Another major trap appears when forming transition-metal cations. Although the 4s orbital fills before the 3d orbitals, 4s electrons are removed first when transition metals form positive ions.
For example, neutral iron is [Ar] 4s² 3d⁶, but Fe²⁺ is [Ar] 3d⁶, not [Ar] 4s² 3d⁴. Always write the neutral atom first, then remove electrons from the 4s orbital before removing them from 3d.