Atomic Structure and Electron Configuration

AP Chemistry· AP Chemistry CED — Atomic Structure and Properties· 14 min read

1. Quantum Numbers and Orbital Structure★★☆☆☆⏱ 4 min

Quantum numbers are a set of four values that describe the unique location and spin of any electron in an atom, consistent with the quantum mechanical model of the atom. Each quantum number narrows down the probability region (orbital) an electron occupies:

Principal quantum number ( $n$ ): Defines the main energy level (shell) and average distance from the nucleus. $n$ is always a positive integer, and larger $n$ corresponds to higher energy and larger orbital size.
Azimuthal (angular momentum) quantum number ( $l$ ): Defines the subshell and shape of the orbital. $l$ can take integer values from $0$ to $n - 1$ , mapped to subshell names: $l = 0 = s$ , $l = 1 = p$ , $l = 2 = d$ , $l = 3 = f$ .
Magnetic quantum number ( $m_{l}$ ): Defines the orientation of the orbital in space, ranging from $- l$ to $+ l$ . The number of orbitals per subshell equals $2 l + 1$ .
Spin quantum number ( $m_{s}$ ): Defines the spin state of the electron, which can only be $+ 1/2$ or $- 1/2$ .

\text{Total electrons per n} = 2n^2

📘 Definition

Orbital

A three-dimensional region around the nucleus where there is a 90% probability of finding a given electron, defined by a unique set of $n$ , $l$ , and $m_{l}$ quantum numbers.

📐 Worked Example

Which of the following sets of quantum numbers $(n, l, m_{l}, m_{s})$ is not allowed for an electron in a ground-state atom? A) $(2, 1, 0, + 1/2)$ B) $(3, 2, - 1, - 1/2)$ C) $(4, 3, + 3, - 1/2)$ D) $(3, 3, - 2, + 1/2)$

1
Recall the first rule for valid quantum numbers: $l$ can only range from $0$ to $n - 1$ , so $l < n$ is always required.
2
Check each option against this rule: A: $l = 1 < 2$ , allowed; B: $l = 2 < 3$ , allowed; C: $l = 3 < 4$ , allowed; D: $l = 3 = n = 3$ , which violates the rule.
3
Confirm: $m_{l} = - 2$ is within the valid range for $l = 3$ , but the violation of $l < n$ makes the entire set invalid.
4
The invalid set is D.

Exam tip:

On AP MCQ, always check if $l \geq n$ first. This is the most common violation tested, so you can eliminate wrong answers in seconds without checking other quantum numbers.

2. Rules for Filling Orbitals★★☆☆☆⏱ 3 min

To write the ground-state electron configuration of an atom, three core rules govern the order electrons fill orbitals:

Aufbau Principle: Electrons fill lower-energy orbitals before higher-energy orbitals. The standard energy order for lighter elements and first-row transition metals is $1 s < 2 s < 2 p < 3 s < 3 p < 4 s < 3 d < 4 p$ .
Pauli Exclusion Principle: No two electrons in the same atom can have identical sets of four quantum numbers. This means each orbital can hold a maximum of two electrons, which must have opposite spins.
Hund's Rule: When filling degenerate (equal-energy) orbitals, electrons occupy each orbital singly with parallel spin before any orbital gets a paired electron. This minimizes electron-electron repulsion.

Two common exceptions to the Aufbau principle for neutral first-row transition metals (frequently tested on the AP exam) are chromium ( $Z = 24$ ) and copper ( $Z = 29$ ). Half-filled ( $d^{5}$ ) and fully filled ( $d^{10}$ ) d subshells have extra stability, so Cr is $[A r] 4 s^{1} 3 d^{5}$ (not $4 s^{2} 3 d^{4}$ ) and Cu is $[A r] 4 s^{1} 3 d^{10}$ (not $4 s^{2} 3 d^{9}$ ).

📐 Worked Example

Draw the ground-state orbital diagram for neutral oxygen ( $Z = 8$ ) and state the number of unpaired electrons.

1
Neutral oxygen has 8 electrons. Fill orbitals in order: 1s holds 2 electrons, 2s holds 2 electrons, leaving 4 electrons for the 2p subshell.
2
Apply the Pauli exclusion principle: 1s and 2s orbitals each have two paired electrons with opposite spin.
3
Apply Hund's rule to the three degenerate 2p orbitals: place one unpaired electron in each orbital first, then pair the fourth electron in one orbital.
4
The final orbital diagram is: $1 s : [↑↓]$ , $2 s : [↑↓]$ , $2 p : [↑↓] [↑] [↑]$ . There are 2 unpaired electrons.

Exam tip:

When drawing orbital diagrams for FRQ, always label each subshell and explicitly show the spin direction of every electron to earn full credit.

3. Electron Configurations for Atoms and Ions★★★☆☆⏱ 4 min

Electron configurations can be written as full configurations (listing all subshells) or condensed (noble gas core) configurations, where the preceding noble gas is placed in brackets to represent inner-shell electrons, and only outer electrons are listed.

The most common point of confusion for students is writing configurations for transition metal cations: the $n s$ electrons are always lost before the $(n - 1) d$ electrons during ionization, even though $n s$ fills before $(n - 1) d$ in neutral atoms. For anions, electrons are added to the lowest available energy subshell, following the same rules as neutral atoms. Valence electrons (the outermost electrons available for bonding) are counted as all electrons with the highest principal quantum number $n$ for main group elements; for transition metals, valence electrons include $n s$ and $(n - 1) d$ electrons.

📐 Worked Example

Write the condensed electron configuration for the $N i^{2 +}$ cation (Ni, $Z = 28$ ).

1
First write the configuration for neutral Ni: $Z = 28$ means 28 electrons. The preceding noble gas is Ar ( $Z = 18$ ), so neutral Ni is $[A r] 4 s^{2} 3 d^{8}$ .
2
Recall the transition metal ionization rule: $4 s$ electrons are lost before $3 d$ electrons. $N i^{2 +}$ has lost 2 electrons total.
3
Remove both electrons from the $4 s$ subshell first, leaving the $3 d$ subshell unchanged.
4
The final condensed configuration for $N i^{2 +}$ is $[A r] 3 d^{8}$ .

✓ Quick check

Test your understanding with this AP-style multiple choice question:

Which of the following is the correct condensed electron configuration for neutral arsenic ( $Z = 33$ )?
- A) $[A r] 4 s^{2} 3 d^{10} 4 p^{3}$
- B) $[A r] 4 s^{2} 4 p^{3}$
- C) $[A r] 3 d^{10} 4 p^{3}$
- D) $[G e] 4 p^{3}$

Exam tip:

After writing any electron configuration, count the total number of electrons to confirm they match the expected number ( $Z$ for neutral, $Z - charge$ for cations, $Z + charge$ for anions). This catches 90% of common counting errors.

4. Common Pitfalls

Wrong move:

Marking the quantum number set $(2, 2, 0, + 1/2)$ as allowed

Why:

Students forget the upper limit of $l$ is $n - 1$ , confusing the maximum $l$ with $n$ .

Correct move:

Always check $l < n$ first when validating quantum number sets.

Wrong move:

Writing the configuration of $T i^{2 +}$ as $[A r] 4 s^{2} 3 d^{0}$

Why:

Students memorize 4s fills before 3d, so incorrectly assume 3d electrons are lost first.

Correct move:

Always remove all $n s$ electrons before removing any $(n - 1) d$ electrons for transition metal cations.

Wrong move:

Drawing the 2p orbital diagram for nitrogen as $[↑↓] [↑] []$

Why:

Students forget Hund's rule and pair electrons early to finish faster.

Correct move:

Always place one unpaired electron into each degenerate orbital before pairing any electrons.

Wrong move:

Writing the electron configuration of neutral copper (Z=29) as $[A r] 4 s^{2} 3 d^{9}$

Why:

Students ignore the extra stability of half-filled and fully filled d subshell exceptions.

Correct move:

Memorize the two AP-tested exceptions: Cr = $[A r] 4 s^{1} 3 d^{5}$ , Cu = $[A r] 4 s^{1} 3 d^{10}$ .

Wrong move:

Counting 18 valence electrons for calcium (Z=20, $[A r] 4 s^{2}$ )

Why:

Students count all electrons instead of only the highest $n$ valence electrons.

Correct move:

For main group elements, only count electrons with the largest principal quantum number $n$ (2 valence electrons for calcium).

Wrong move:

Assigning $l = 1$ to an s subshell

Why:

Students mix up the mapping of $l$ values to subshell names.

Correct move:

Memorize the mapping: $l = 0 = s, l = 1 = p, l = 2 = d, l = 3 = f$ .

5. Quick Reference Cheatsheet

Category	Formula/Rule	Notes
Maximum electrons per principal shell $n$	$2 n^{2}$	Counts all electrons in all subshells with principal quantum number $n$
Azimuthal quantum number range	$0 \leq l \leq n - 1$	Mapping: $l = 0 = s, l = 1 = p, l = 2 = d, l = 3 = f$ ; check this first for valid quantum numbers
Magnetic quantum number range	$- l \leq m_{l} \leq + l$	Number of orbitals per subshell = $2 l + 1$
Pauli Exclusion Principle	2 electrons per orbital, opposite spin	No two electrons share the same full set of four quantum numbers
Hund's Rule	Fill degenerate orbitals with parallel spin first	Minimizes electron-electron repulsion for ground state configurations
Aufbau filling order	$1 s < 2 s < 2 p < 3 s < 3 p < 4 s < 3 d < 4 p$	Exceptions: Cr = $[A r] 4 s^{1} 3 d^{5}$ , Cu = $[A r] 4 s^{1} 3 d^{10}$
Transition metal ionization	Lose $n s$ electrons before $(n - 1) d$ electrons	The most tested rule after basic configuration writing
Valence electron count (main group)	Number of electrons with highest $n$	Matches the group number for main group elements
Electron count for charged species	Cations: $Z - charge$ ; Anions: $Z + charge$	Always count after writing a configuration to catch errors

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2023 · MCQ
Identify invalid quantum number set
2022 · FRQ
Write ion electron configuration
2021 · MCQ
Paramagnetism and unpaired electrons

Going deeper

unit overviewAP Chemistry Unit 1 Overview

What's Next

This topic is the absolute foundation for all subsequent topics in AP Chemistry, because electron arrangement directly determines every chemical property of an element, from its reactivity to its bonding behavior. Next, you will apply the rules of electron configuration to understand periodic trends, including atomic radius, ionization energy, and electron affinity, which are heavily tested on the AP exam. Without correctly writing electron configurations and counting valence electrons, you cannot explain why these trends exist or predict the properties of unfamiliar elements. This topic also feeds directly into the study of chemical bonding, molecular geometry, and intermolecular forces later in the course.