# Electron orbitals and configuration

> CIE A-Level Chemistry · 9701
> Source: https://www.owlsprep.com/study/cie-9701-u2-electron-orbitals-and-configuration/

This module explains quantum electron structure, including shells, subshells, orbitals, and how to write full/abbreviated electron configurations for atoms and ions up to Z=36. It covers CIE-specific common exceptions and exam traps.

**Prerequisites:** [Basic atomic structure (protons, neutrons, electrons)](https://www.owlsprep.com/study/cie-9701-u1-structure-of-the-atom/)

## Learning objectives

- Describe the structure of electron shells, subshells and orbitals
- Write electron configurations for atoms and ions up to Z = 36
- State and apply the Aufbau principle, Pauli exclusion principle and Hund's rule
- Explain the common exceptions to the standard Aufbau filling order

## Quantum Structure: Shells, Subshells and Orbitals

**Atomic Orbital** — A 3-dimensional region of space around the nucleus where there is a 95% probability of finding an electron with a specific fixed energy.

*Example:* A 1s orbital is a spherical region closest to the nucleus.

In the modern quantum model, electrons are arranged in principal shells (energy levels) labelled by the principal quantum number $n=1,2,3...$. Lower $n$ corresponds to lower energy and smaller distance from the nucleus. Each principal shell splits into subshells, which each contain a fixed number of orbitals.

| Subshell letter | Number of orbitals | Maximum electrons | First appears at n= |
| --- | --- | --- | --- |
| s | 1 | 2 | 1 |
| p | 3 | 6 | 2 |
| d | 5 | 10 | 3 |
| f | 7 | 14 | 4 |

> **info**
>
> The Pauli Exclusion Principle states that no more than 2 electrons can occupy the same orbital, and they must have opposite spins.

**Worked example:** Calculate the maximum number of electrons that can occupy the $n=3$ principal shell.

1. Identify all subshells present in $n=3$:
2. $$3s, 3p, 3d$$
3. Find the maximum electrons per subshell:
4. 3s = 2, 3p = 6, 3d = 10
5. Sum the total number of electrons:
6. $$2 + 6 + 10 = 18$$

*Calculator:* forbidden

## Rules for Filling Orbitals

**Aufbau Principle** — Electrons fill orbitals in order of increasing energy: lower energy orbitals are filled completely before higher energy orbitals are occupied.

1. **Pauli Exclusion Principle**: Each orbital holds a maximum of 2 electrons with opposite spins.
2. **Hund's Rule**: When filling degenerate (equal energy) orbitals (e.g. the three 2p orbitals), electrons fill each orbital singly first with parallel spins, before pairing up. This minimizes electron-electron repulsion.

For elements up to Z=36 (krypton), the order of filling orbitals by increasing energy is:

$$1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p$$

> **mnemonic**
>
> Use the diagonal rule to remember filling order: draw diagonals from top-right to bottom-left through this grid to get the correct order: <br> 1s <br> 2s 2p <br> 3s 3p 3d <br> 4s 4p 4d 4f

**Worked example:** Draw the orbital filling diagram for neutral nitrogen (Z=7).

1. Total number of electrons = atomic number = 7.
2. Fill lowest energy orbitals first: 1s holds 2 paired electrons, 2s holds 2 paired electrons, leaving $7 - 4 = 3$ electrons for 2p.
3. Apply Hund's rule: 2p has 3 degenerate orbitals, so add one unpaired electron to each orbital with parallel spin.

> **Exam tip:** CIE examiners regularly test application of Hund's rule, so always check for single occupation before pairing electrons in orbitals.

*Calculator:* forbidden

## Writing Electron Configurations

Electron configurations can be written in full, or abbreviated using noble gas core shorthand. For CIE exams, you are expected to write configurations for all atoms and ions up to Z=36.

> **note**
>
> Convention: after 3d starts filling, 4s becomes higher energy than 3d. We always write configurations sorted by principal quantum number $n$, so 3d is written before 4s, even though 4s fills first.

**Worked example:** Write the full and abbreviated electron configuration for neutral iron (Z=26).

1. Total electrons = 26, fill in the energy order: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6$
2. Rearrange to conventional order sorted by $n$:
3. $$1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2$$
4. For abbreviated form, replace the core electrons with the symbol of the previous noble gas (argon, Z=18):
5. $$[Ar] 3d^6 4s^2$$

**Check your understanding**

Test your understanding of transition metal ion configurations:

1. What is the correct electron configuration of $Fe^{2+}$?

   - [Ar] 3d^6
   - [Ar] 3d^4 4s^2
   - [Ar] 3d^6 4s^2
   - [Ar] 3d^8

   *Why:* 4s electrons are always lost before 3d electrons when transition metals form positive ions. Neutral Fe has 2 electrons in 4s, so both are lost to form Fe²⁺.

*Calculator:* forbidden

## Exceptions to the Aufbau Principle

For CIE 9701, you only need to remember two common exceptions to the standard filling order for first-row transition metals: chromium (Z=24) and copper (Z=29). These exceptions occur because a half-filled (5 electrons) or fully filled (10 electrons) 3d subshell has extra stability from symmetry and exchange energy.

**Worked example:** Write the correct electron configuration for copper (Z=29) and explain why it is an exception.

1. Standard filling would give the configuration $[Ar] 3d^9 4s^2$.
2. A fully filled 3d subshell (10 electrons) has extra stability compared to a $3d^9$ arrangement.
3. One electron from 4s is promoted to 3d to achieve this stable configuration.
4. The correct configuration is:
5. $$[Ar] 3d^{10} 4s^1$$

> **Exam tip:** CIE almost always tests these two exceptions, so memorise: Cr = [Ar] 3d⁵ 4s¹, Cu = [Ar] 3d¹⁰ 4s¹.

*Calculator:* forbidden

## Common pitfalls

- **Wrong:** Claims 4s is always written before 3d because it fills first
  - Why it fails: Confuses the order of filling with the conventional writing order sorted by principal quantum number n
  - Correct: 4s fills before 3d, but 3d has lower n so it is written before 4s in the final configuration
- **Wrong:** Writes chromium as [Ar] 3d⁴ 4s²
  - Why it fails: Forgets the extra stability of a half-filled 3d subshell, the key exception tested by CIE
  - Correct: The correct configuration for chromium is [Ar] 3d⁵ 4s¹
- **Wrong:** Removes 3d electrons first when forming transition metal ions
  - Why it fails: Assumes 4s is always lower energy than 3d, which is only true before 3d starts filling
  - Correct: After 3d fills, 4s is higher energy, so 4s electrons are always lost first when ionising
- **Wrong:** Defines an orbital as a fixed path electrons follow
  - Why it fails: Confuses the old Bohr model with the modern quantum model
  - Correct: An orbital is a region of space where there is ~95% probability of finding an electron of a given energy
- **Wrong:** Pairs electrons in degenerate orbitals before all have one electron
  - Why it fails: Violates Hund's rule of maximum multiplicity
  - Correct: Fill all degenerate orbitals with one parallel-spin electron before pairing any electrons

## Cheatsheet

| Concept | Key CIE Exam Fact |
| --- | --- |
| Filling order | 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p |
| Writing order | Sort by n: 1s < 2s < 2p < 3s < 3p < 3d < 4s < 4p |
| Max electrons per subshell | s=2, p=6, d=10, f=14 |
| Common exceptions | Cr: [Ar] 3d⁵ 4s¹; Cu: [Ar] 3d¹⁰ 4s¹ |
| Ion electron loss | 4s electrons lost before 3d electrons |
| Orbital definition | Region of 95% probability of finding an electron |

## What's next

Electron configuration is the foundation of all modern chemistry, explaining periodic trends, bonding patterns, chemical reactivity and the unique properties of transition metals. This sub-topic is assessed in almost every CIE 9701 paper, appearing in both multiple choice and structured questions, so mastering the rules and exceptions is critical for achieving a high grade. Understanding how electron arrangement relates to ionisation energy will help you tackle periodic trend questions later in the course.

- [Ionisation energies](https://www.owlsprep.com/study/cie-9701-u2-ionisation-energies/)
- [Chemical bonding](https://www.owlsprep.com/study/cie-9701-u3-overview/)
- [Covalent and coordinate bonding](https://www.owlsprep.com/study/cie-9701-u3-covalent-and-coordinate-bonding/)

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