Electronic Structure

This lesson covers the electronic structure (electron configuration) of atoms, as required by AQA GCSE Chemistry specification (5.1.1). The arrangement of electrons in an atom determines how that element reacts, what type of bonds it forms, and where it sits in the periodic table. Understanding electronic structure is essential for explaining reactivity, bonding, and the properties of elements.

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Electron Shells (Energy Levels)

Electrons occupy shells (also called energy levels) around the nucleus. Each shell can hold a maximum number of electrons:

Shell Number	Maximum Electrons	Distance from Nucleus
1st shell (innermost)	2	Closest
2nd shell	8	Further
3rd shell	8 (at GCSE level)	Furthest (for first 20 elements)

Rules for Filling Electron Shells

1. Electrons always fill the lowest energy level (closest to the nucleus) first.
2. When a shell is full, electrons begin to fill the next shell.
3. The first shell fills first (max 2), then the second shell (max 8), then the third shell (max 8 for GCSE).
4. For the first 20 elements, this simple rule works. Beyond element 20 (calcium), the filling pattern becomes more complex (covered at A-level).

Exam Tip: You must be able to write the electronic structure for the first 20 elements. The filling order is straightforward: 1st shell (max 2), 2nd shell (max 8), 3rd shell (max 8). Always fill each shell completely before starting the next one.

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Electronic Configurations of the First 20 Elements

The electronic configuration is written as numbers separated by dots or commas, showing how many electrons are in each shell.

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

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Drawing Electron Shell Diagrams

You may be asked to draw electron shell diagrams in the exam. Follow these steps:

1. Draw a small circle in the centre to represent the nucleus — write the element symbol inside.
2. Draw concentric circles around it to represent each occupied shell.
3. Place dots or crosses on each circle to show the electrons.
4. Label the number of electrons in each shell.

Example: Sodium (2, 8, 1)

• Nucleus: 11 protons, 12 neutrons.
• 1st shell: 2 electrons.
• 2nd shell: 8 electrons.
• 3rd shell: 1 electron.

Exam Tip: When drawing electron shell diagrams, always make sure the number of electrons equals the atomic number (for neutral atoms). Count your electrons carefully — a common mistake is to draw the wrong number of electrons on a shell.

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Electronic Structure and the Periodic Table

The electronic configuration of an element directly determines its position in the periodic table:

Feature	Determined by
Group number	Number of electrons in the outermost shell (valence electrons)
Period number	Number of occupied electron shells

For example:

• Sodium (2, 8, 1) — 1 electron in the outer shell = Group 1; 3 occupied shells = Period 3.
• Chlorine (2, 8, 7) — 7 electrons in the outer shell = Group 7; 3 occupied shells = Period 3.
• Neon (2, 8) — 8 electrons in the outer shell = Group 0 (noble gases, full outer shell); 2 occupied shells = Period 2.

graph LR
    A["Electronic Configuration"] --> B["Number of outer electrons = Group number"]
    A --> C["Number of shells = Period number"]
    B --> D["Group 1: 1 outer electron"]
    B --> E["Group 7: 7 outer electrons"]
    B --> F["Group 0: full outer shell"]
    C --> G["Period 1: 1 shell"]
    C --> H["Period 2: 2 shells"]
    C --> I["Period 3: 3 shells"]

    style A fill:#2c3e50,color:#fff
    style B fill:#e74c3c,color:#fff
    style C fill:#3498db,color:#fff

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Outer Electrons and Reactivity

The electrons in the outermost shell (valence electrons) are the most important electrons in chemistry because they are the ones involved in chemical bonding and reactions.

Why Outer Electrons Matter

• Atoms react to achieve a stable electron configuration — usually a full outer shell.
• Elements with full outer shells (noble gases) are very stable and unreactive.
• Elements with one or two outer electrons (metals) tend to lose electrons to form positive ions.
• Elements with six or seven outer electrons (non-metals) tend to gain electrons to form negative ions.
• Elements with four outer electrons (like carbon and silicon) tend to share electrons in covalent bonds.

Group	Outer Electrons	Tendency	Ion Formed
1 (Alkali metals)	1	Lose 1 electron	+1 ion (e.g. Na+)
2 (Alkaline earth metals)	2	Lose 2 electrons	+2 ion (e.g. Mg2+)
6	6	Gain 2 electrons	-2 ion (e.g. O2-)
7 (Halogens)	7	Gain 1 electron	-1 ion (e.g. Cl-)
0 (Noble gases)	Full shell (2 or 8)	No tendency to lose or gain	No ion — already stable

Exam Tip: Atoms always react to achieve the electronic structure of the nearest noble gas. Metals lose electrons to look like the previous noble gas; non-metals gain electrons to look like the next noble gas. This concept of achieving a full outer shell is central to understanding bonding.

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Ions and Electronic Structure

When atoms lose or gain electrons, they become ions — charged particles. The electronic structure of an ion is different from that of the neutral atom.

Example: Sodium atom to sodium ion

• Sodium atom (Na): 2, 8, 1 — has 11 protons and 11 electrons (neutral).
• Sodium ion (Na+): 2, 8 — has 11 protons and 10 electrons (lost 1 electron, now positive).
• The sodium ion has the same electronic structure as neon (the nearest noble gas).