Electronic Configuration

This lesson covers how electrons are arranged in atoms, as required by AQA GCSE Combined Science Trilogy (8464, Chemistry 4.1.1). Electronic configuration determines how an element reacts and where it sits in the periodic table. Understanding this topic is essential for explaining bonding, reactivity trends and the periodic table.

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Energy Levels (Electron Shells)

Electrons occupy energy levels (also called shells) around the nucleus. The shells are filled from the innermost shell (closest to the nucleus, lowest energy) outwards.

Energy Level (Shell)	Maximum Electrons
1st shell	2
2nd shell	8
3rd shell	8 (at GCSE level)

Exam Tip: At GCSE, the 3rd shell holds a maximum of 8 electrons. In reality it can hold 18, but the additional electrons fill 3d sub-shells, which is an A-level topic. AQA will only test up to the first 20 elements.

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Rules for Filling Electron Shells

1. Electrons always fill the lowest available energy level first.
2. Once a shell is full, electrons begin filling the next shell.
3. For the first 20 elements, this means: fill the 1st shell (2), then the 2nd shell (8), then the 3rd shell (8), then the 4th shell starts.

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Electronic Configurations of the First 20 Elements

Element	Symbol	Atomic Number	Electronic Configuration
Hydrogen	H	1	$1$1
Helium	He	2	$2$2
Lithium	Li	3	$2, 1$2,1
Beryllium	Be	4	$2, 2$2,2
Boron	B	5	$2, 3$2,3
Carbon	C	6	$2, 4$2,4
Nitrogen	N	7	$2, 5$2,5
Oxygen	O	8	$2, 6$2,6
Fluorine	F	9	$2, 7$2,7
Neon	Ne	10	$2, 8$2,8
Sodium	Na	11	$2, 8, 1$2,8,1
Magnesium	Mg	12	$2, 8, 2$2,8,2
Aluminium	Al	13	$2, 8, 3$2,8,3
Silicon	Si	14	$2, 8, 4$2,8,4
Phosphorus	P	15	$2, 8, 5$2,8,5
Sulfur	S	16	$2, 8, 6$2,8,6
Chlorine	Cl	17	$2, 8, 7$2,8,7
Argon	Ar	18	$2, 8, 8$2,8,8
Potassium	K	19	$2, 8, 8, 1$2,8,8,1
Calcium	Ca	20	$2, 8, 8, 2$2,8,8,2

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Electronic Configuration and the Periodic Table

The electronic configuration of an element is directly linked to its position in the periodic table:

Feature	What It Tells You
Group number	The number of electrons in the outer shell
Period number	The number of occupied electron shells

Examples

• Sodium ($2, 8, 1$2,8,1) — Group 1 (1 electron in outer shell), Period 3 (3 shells occupied).
• Chlorine ($2, 8, 7$2,8,7) — Group 7 (7 electrons in outer shell), Period 3 (3 shells).
• Calcium ($2, 8, 8, 2$2,8,8,2) — Group 2 (2 electrons in outer shell), Period 4 (4 shells).

Exam Tip (AQA 8464): If you are given an electronic configuration, you can immediately identify the group (outer shell electrons) and period (number of shells). This is a very common short-answer question.

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Why Electronic Configuration Matters

The outer shell electrons (also called valence electrons) determine:

1. Chemical reactivity — how easily an atom gains, loses or shares electrons.
2. Which group the element belongs to.
3. The type of bonding the element forms (ionic or covalent).
4. Similarity of properties within a group — elements in the same group have the same number of outer electrons, so they react in similar ways.

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Worked Example

Question: An element has the electronic configuration $2, 8, 6$2,8,6. Identify its group, period and name.

Answer:

• Outer shell has 6 electrons → Group 6
• 3 shells occupied → Period 3
• Atomic number = 2 + 8 + 6 = 16 → this is sulfur (S)

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Drawing Electronic Configurations

When drawing electronic configurations:

1. Draw the nucleus in the centre (you can write the atomic number or the number of protons/neutrons).
2. Draw concentric circles for each energy level.
3. Place dots or crosses to represent electrons in each shell.
4. Fill shells from the inside out.

For example, sodium ($2, 8, 1$2,8,1):

• 1st shell: 2 electrons
• 2nd shell: 8 electrons
• 3rd shell: 1 electron

graph TD
    A["Nucleus: 11p, 12n"] --> B["1st Shell: 2 electrons"]
    B --> C["2nd Shell: 8 electrons"]
    C --> D["3rd Shell: 1 electron"]
    D --> E["Configuration: 2, 8, 1"]
    E --> F["Group 1, Period 3"]

    style A fill:#e74c3c,color:#fff
    style B fill:#3498db,color:#fff
    style C fill:#3498db,color:#fff
    style D fill:#3498db,color:#fff
    style E fill:#2ecc71,color:#fff
    style F fill:#9b59b6,color:#fff

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Stable Electronic Configurations

Atoms with full outer shells have a stable electronic configuration and are very unreactive. This is why the noble gases (Group 0) are so stable:

• Helium: $2$2 (full 1st shell)
• Neon: $2, 8$2,8 (full 2nd shell)
• Argon: $2, 8, 8$2,8,8 (full 3rd shell at GCSE level)

Other elements react in order to achieve a full outer shell — by losing, gaining or sharing electrons.

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Common Mistakes

Mistake	Correction
Putting more than 2 electrons in the 1st shell	The 1st shell holds a maximum of 2 electrons
Putting more than 8 in the 2nd or 3rd shell (at GCSE)	At GCSE, shells 2 and 3 hold a maximum of 8
Confusing the group number with the period number	Group = outer shell electrons; Period = number of shells
Not knowing that potassium is $2, 8, 8, 1$2,8,8,1	After the 3rd shell fills to 8, the 4th shell starts
Saying helium is in Group 2 because it has 2 outer electrons	Helium is in Group 0 — its outer shell (1st shell) is full

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Summary

• Electrons are arranged in energy levels (shells): 1st holds 2, 2nd holds 8, 3rd holds 8 (at GCSE).
• Electrons fill the lowest energy level first.
• The group number equals the number of outer shell electrons.
• The period number equals the number of occupied shells.
• Atoms with full outer shells are stable and unreactive.
• Elements in the same group have similar properties because they have the same number of outer shell electrons.