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This lesson covers the arrangement of electrons in atoms, including shells, sub-shells, and orbitals, the rules governing electron filling, and the important exceptions at A-Level. Electron configuration underpins the whole of chemistry — from bonding to reactivity to periodicity.
Spec mapping (AQA 7405): This lesson maps to §3.1.1.3 (electron configuration, including shells, sub-shells, orbitals, and the s-, p-, d- and f-blocks of the periodic table). The half-filled and fully-filled-stability exceptions for chromium and copper are explicitly listed in the spec. The link to ionisation energies (§3.1.1.3 continued) is examined in the next lesson. Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: Writing electron configurations for atoms and ions is straight AO1 / AO2 territory and appears on every paper. Explaining the Cr / Cu exceptions in terms of half-filled / fully-filled stability is a recurring AO3 item, often worth 3 marks. Predicting block / group / period from a given configuration is AO2.
Electrons are arranged in shells (principal energy levels), numbered n = 1, 2, 3, 4, etc. The shell number n determines:
| Shell (n) | Sub-shells | Maximum electrons (2n²) |
|---|---|---|
| 1 | 1s | 2 |
| 2 | 2s, 2p | 8 |
| 3 | 3s, 3p, 3d | 18 |
| 4 | 4s, 4p, 4d, 4f | 32 |
Each shell is divided into sub-shells labelled s, p, d, and f. These sub-shells have different energies (within the same shell, s < p < d < f).
| Sub-shell | Number of orbitals | Maximum electrons |
|---|---|---|
| s | 1 | 2 |
| p | 3 | 6 |
| d | 5 | 10 |
| f | 7 | 14 |
An orbital is a region of space around the nucleus where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons, which must have opposite spins (spin-up ↑ and spin-down ↓).
Shapes of orbitals:
Key Point: All orbitals in the same sub-shell have the same energy — they are described as degenerate. For example, the three 2p orbitals are degenerate.
Three rules govern how electrons fill orbitals:
Electrons fill orbitals in order of increasing energy. The filling order is:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p
Note that 4s fills before 3d because 4s has a lower energy in neutral atoms. This is the most important exception to the simple numerical ordering.
Within a sub-shell, electrons occupy orbitals singly before pairing up. All singly occupied orbitals have electrons with the same spin (parallel spins). This minimises electron-electron repulsion.
For example, in nitrogen (7 electrons): 1s² 2s² 2p³ The three 2p electrons occupy the three 2p orbitals singly:
2p: [↑] [↑] [↑] — correct (Hund's rule) 2p: [↑↓] [↑] [ ] — incorrect (electrons should not pair before all orbitals have one)
No two electrons in an atom can have the same set of four quantum numbers. In practice, this means each orbital can hold a maximum of two electrons, and they must have opposite spins.
The full electron configuration lists all sub-shells with the number of electrons as a superscript.
Examples:
| Element | Z | Electron configuration |
|---|---|---|
| Hydrogen | 1 | 1s¹ |
| Helium | 2 | 1s² |
| Lithium | 3 | 1s² 2s¹ |
| Carbon | 6 | 1s² 2s² 2p² |
| Nitrogen | 7 | 1s² 2s² 2p³ |
| Oxygen | 8 | 1s² 2s² 2p⁴ |
| Neon | 10 | 1s² 2s² 2p⁶ |
| Sodium | 11 | 1s² 2s² 2p⁶ 3s¹ |
| Argon | 18 | 1s² 2s² 2p⁶ 3s² 3p⁶ |
| Potassium | 19 | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ |
| Calcium | 20 | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² |
| Scandium | 21 | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s² |
| Titanium | 22 | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s² |
| Iron | 26 | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² |
Exam Tip: When writing electron configurations, always write 3d before 4s (even though 4s fills first), because you are listing sub-shells in order of principal quantum number. For example, scandium is written as ...3d¹ 4s², not ...4s² 3d¹.
For elements beyond the first period, you can abbreviate the inner electron configuration using the noble gas core:
Two transition metals have electron configurations that deviate from the expected filling order:
Expected: [Ar] 3d⁴ 4s² Actual: [Ar] 3d⁵ 4s¹
A half-filled 3d sub-shell (3d⁵) has extra stability due to the symmetrical distribution of electrons and the exchange energy associated with five electrons of the same spin. One electron is promoted from 4s to 3d.
Expected: [Ar] 3d⁹ 4s² Actual: [Ar] 3d¹⁰ 4s¹
A fully-filled 3d sub-shell (3d¹⁰) has extra stability. One electron is promoted from 4s to 3d to achieve this.
Common Misconception: Students often forget these exceptions. The AQA specification explicitly requires you to know the electron configurations of chromium and copper. Remember: Cr is [Ar] 3d⁵ 4s¹ and Cu is [Ar] 3d¹⁰ 4s¹.
When atoms lose electrons to form positive ions, the electrons are removed from the highest energy sub-shell first. For transition metals, this means 4s electrons are removed before 3d electrons.
Examples:
| Ion | Configuration |
|---|---|
| Na⁺ | 1s² 2s² 2p⁶ |
| Mg²⁺ | 1s² 2s² 2p⁶ |
| Fe²⁺ | [Ar] 3d⁶ (lost two 4s electrons) |
| Fe³⁺ | [Ar] 3d⁵ (lost two 4s and one 3d electron) |
| Cu²⁺ | [Ar] 3d⁹ (lost one 4s and one 3d electron from [Ar] 3d¹⁰ 4s¹) |
Key Point: 4s electrons are removed first when forming ions, even though 4s fills before 3d. This is because once 3d is occupied, 4s becomes higher in energy than 3d.
When atoms gain electrons to form negative ions, electrons are added to the lowest available energy sub-shell.
| Ion | Configuration |
|---|---|
| F⁻ | 1s² 2s² 2p⁶ |
| O²⁻ | 1s² 2s² 2p⁶ |
| Cl⁻ | 1s² 2s² 2p⁶ 3s² 3p⁶ |
An electron-in-boxes (or orbital) diagram represents each orbital as a box and each electron as an arrow. Arrows pointing up (↑) and down (↓) represent the two possible spins.
Configuration: 1s² 2s² 2p⁴
1s: [↑↓] 2s: [↑↓] 2p: [↑↓] [↑] [↑]
By Hund's rule, the first three 2p electrons go in singly, then the fourth pairs up in the first 2p orbital.
Configuration: [Ar] 3d⁶ 4s²
3d: [↑↓] [↑] [↑] [↑] [↑] 4s: [↑↓]
Five 3d electrons fill singly first (Hund's rule), then the sixth pairs up.
Configuration: [Ar] 3d⁵ 4s¹
3d: [↑] [↑] [↑] [↑] [↑] 4s: [↑]
All 3d orbitals are singly occupied — maximum exchange energy and stability.
The periodic table is structured according to electron filling:
| Block | Sub-shell being filled | Groups |
|---|---|---|
| s-block | s sub-shell | 1–2 (+ He) |
| p-block | p sub-shell | 13–18 |
| d-block | d sub-shell | 3–12 |
| f-block | f sub-shell | Lanthanides and actinides |
The period number corresponds to the outermost shell being filled. The group number is related to the number of outer-shell electrons (for main group elements).
An element has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁴. Identify the element and predict its group and period.
Exam Tip: You may be given an electron configuration and asked to identify the element, its position in the periodic table, or whether it is in its ground state. An element is in its ground state when electrons fill orbitals according to the Aufbau principle. An excited state has one or more electrons promoted to a higher energy level.
Electron configuration is one of the most-connected topics in A-Level Chemistry:
Question 1. [12 marks total]
(a) Write the full electron configuration of an atom of iron (Z = 26). [1 mark]
(b) Write the full electron configuration of an Fe³⁺ ion. Explain which electrons are removed first when forming this ion. [3 marks]
(c) Explain why the electron configuration of chromium (Z = 24) is [Ar] 3d⁵ 4s¹ rather than the expected [Ar] 3d⁴ 4s². [3 marks]
(d) State Hund's rule. Using nitrogen as an example, draw an electron-in-boxes diagram for the 2p sub-shell and explain how it illustrates the rule. [3 marks]
(e) An element X has the electron configuration [Ar] 3d¹⁰ 4s² 4p². State the block, period and group of element X, and identify the element. [2 marks]
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