What you'll learn
This revision guide covers electronic structure, a fundamental topic in AQA GCSE Chemistry that explains how electrons are arranged in atoms. You'll learn how to represent electron arrangements, understand the link between electronic structure and the periodic table, and explain chemical reactivity in terms of electron configuration. This knowledge underpins bonding, the periodic table, and much of your Chemistry GCSE.
Key terms and definitions
Electron shell — an energy level within an atom where electrons orbit the nucleus. Also called an electron energy level.
Electronic structure — the arrangement of electrons in shells around the nucleus of an atom, written as a series of numbers (e.g., 2,8,1 for sodium).
Valence electrons — electrons in the outermost shell of an atom, which determine how the element reacts chemically.
Noble gas configuration — a stable electronic structure with a full outer shell, characteristic of Group 0 elements (helium, neon, argon, etc.).
Electron shell diagram — a visual representation showing electrons arranged in concentric circles around the nucleus.
Ground state — the lowest energy state of an atom, with electrons occupying the lowest available shells.
Octet rule — the tendency of atoms to gain, lose, or share electrons to achieve eight electrons in their outer shell (or two for the first shell).
Atomic number — the number of protons in the nucleus of an atom, which equals the number of electrons in a neutral atom.
Core concepts
Electron shells and their capacities
Electrons occupy shells (energy levels) around the nucleus. Each shell can hold a maximum number of electrons:
- First shell (closest to nucleus): maximum 2 electrons
- Second shell: maximum 8 electrons
- Third shell: maximum 8 electrons (at GCSE level)
- Fourth shell: maximum 2 electrons for elements 1-20 only
Electrons fill shells starting from the innermost shell outwards. The first shell must be full before electrons occupy the second shell, and so on. This systematic filling pattern is called the ground state configuration.
For elements with atomic numbers 1-20 (hydrogen to calcium), you need to know how to work out and represent electronic structures. Beyond calcium, the pattern becomes more complex and is not required at GCSE.
The general rule for elements 1-20:
- Find the atomic number (number of protons = number of electrons)
- Place 2 electrons in the first shell
- Place up to 8 electrons in the second shell
- Place up to 8 electrons in the third shell
- Any remaining electrons go in the fourth shell
Writing electronic structures
Electronic structures are written as a series of numbers separated by commas, showing the number of electrons in each shell from innermost to outermost.
Examples for common elements:
- Hydrogen (H, atomic number 1): 1
- Helium (He, atomic number 2): 2
- Carbon (C, atomic number 6): 2,4
- Oxygen (O, atomic number 8): 2,6
- Sodium (Na, atomic number 11): 2,8,1
- Chlorine (Cl, atomic number 17): 2,8,7
- Calcium (Ca, atomic number 20): 2,8,8,2
Notice that the sum of all numbers equals the atomic number. This is a useful check when writing electronic structures.
For ions, the number of electrons differs from the atomic number:
- Positive ions (cations) have lost electrons: Na⁺ has 10 electrons, so electronic structure is 2,8
- Negative ions (anions) have gained electrons: Cl⁻ has 18 electrons, so electronic structure is 2,8,8
Drawing electron shell diagrams
Electron shell diagrams provide a visual representation of electronic structure. Follow these conventions:
Drawing method:
- Draw the nucleus as a small circle in the centre, labelled with the element symbol or containing protons and neutrons
- Draw concentric circles around the nucleus representing shells
- Mark electrons as crosses (×), dots (•), or small circles on each shell
- Distribute electrons correctly according to shell capacity
- Space electrons evenly around each shell for clarity
Key points for accurate diagrams:
- Shells should be evenly spaced and roughly circular
- Electrons are positioned on the line of the circle (not inside or outside)
- For the first few electrons in a shell, space them out evenly
- Label or make clear which element is shown
This representation helps visualise the electronic structure and is particularly useful when explaining bonding and reactivity.
Electronic structure and the periodic table
The periodic table is organized according to electronic structure. This organization explains the table's structure and predicts chemical properties.
Group number and outer electrons:
For Groups 1, 2, and 13-18, the group number relates directly to the number of outer shell electrons:
- Group 1 elements have 1 electron in the outer shell (e.g., Li: 2,1; Na: 2,8,1)
- Group 2 elements have 2 electrons in the outer shell (e.g., Mg: 2,8,2; Ca: 2,8,8,2)
- Group 13 elements have 3 electrons in the outer shell (e.g., Al: 2,8,3)
- Group 14 elements have 4 electrons in the outer shell (e.g., C: 2,4; Si: 2,8,4)
- Group 15 elements have 5 electrons in the outer shell (e.g., N: 2,5; P: 2,8,5)
- Group 16 elements have 6 electrons in the outer shell (e.g., O: 2,6; S: 2,8,6)
- Group 17 elements have 7 electrons in the outer shell (e.g., F: 2,7; Cl: 2,8,7)
- Group 18 (Group 0) elements have 8 electrons in the outer shell, except helium which has 2 (e.g., Ne: 2,8; Ar: 2,8,8)
Period number and shells:
The period number (row) indicates the number of occupied electron shells:
- Period 1 elements (H, He) have electrons in 1 shell
- Period 2 elements (Li to Ne) have electrons in 2 shells
- Period 3 elements (Na to Ar) have electrons in 3 shells
- Period 4 elements (K, Ca at GCSE) have electrons in 4 shells
This pattern explains why the periodic table has its characteristic shape and why elements in the same group have similar chemical properties.
Electronic structure and reactivity
Chemical reactions involve the loss, gain, or sharing of outer shell electrons. Elements react to achieve a stable electronic structure, typically a full outer shell like the noble gases.
Noble gases (Group 0/18):
- Already have full outer shells (8 electrons, or 2 for helium)
- Very unreactive (stable configuration)
- Do not readily form ions or bonds
- Provide the "target" structure that other elements try to achieve
Group 1 metals (alkali metals):
- Have 1 electron in the outer shell
- React by losing this electron to form +1 ions
- This gives them the electronic structure of the previous noble gas
- Example: Na (2,8,1) → Na⁺ (2,8), same as neon
Group 7 non-metals (halogens):
- Have 7 electrons in the outer shell
- React by gaining 1 electron to form -1 ions
- This gives them the electronic structure of the next noble gas
- Example: Cl (2,8,7) → Cl⁻ (2,8,8), same as argon
Group 2 metals:
- Have 2 electrons in the outer shell
- React by losing both electrons to form +2 ions
- Example: Mg (2,8,2) → Mg²⁺ (2,8), same as neon
Group 6 non-metals:
- Have 6 electrons in the outer shell
- React by gaining 2 electrons to form -2 ions
- Example: O (2,6) → O²⁻ (2,8), same as neon
Elements in the same group react similarly because they have the same number of outer shell electrons. This fundamental principle explains patterns in reactivity across the periodic table.
Limitations and extensions
At GCSE level, several simplifications apply:
Third shell capacity: The third shell can actually hold up to 18 electrons, but at GCSE you treat it as holding a maximum of 8 electrons for the first 20 elements. This is sufficient for understanding chemical properties of these elements.
Transition metals: Transition metals (Groups 3-12) have more complex electronic structures involving d-orbitals. These are not covered in detail at GCSE. You only need to know that transition metals are in the central block of the periodic table and can form ions with different charges.
Sub-shells: Electron shells are actually divided into sub-shells (s, p, d, f). At GCSE, you don't need to know about sub-shells. The simplified shell model is sufficient for explaining reactivity and bonding.
Worked examples
Example 1: Working out electronic structure
Question: Phosphorus has an atomic number of 15. Write the electronic structure of: (a) a phosphorus atom [1 mark] (b) a phosphide ion, P³⁻ [1 mark]
Solution:
(a) Phosphorus atom (15 electrons):
- First shell: 2 electrons
- Second shell: 8 electrons
- Third shell: 5 electrons (15 - 2 - 8 = 5)
- Electronic structure: 2,8,5 ✓
(b) Phosphide ion P³⁻ (15 + 3 = 18 electrons):
- First shell: 2 electrons
- Second shell: 8 electrons
- Third shell: 8 electrons
- Electronic structure: 2,8,8 ✓
Mark scheme notes: Award 1 mark for each correct answer. Common error: forgetting to add electrons for negative ions or subtract for positive ions.
Example 2: Explaining reactivity using electronic structure
Question: Explain, in terms of electronic structure, why sodium and potassium have similar chemical properties. [3 marks]
Solution:
Sodium has electronic structure 2,8,1 and potassium has electronic structure 2,8,8,1 ✓
Both elements have 1 electron in their outer shell ✓
Elements with the same number of outer shell electrons react in similar ways / both lose one electron to form +1 ions ✓
Mark scheme notes:
- 1 mark for stating or implying both have 1 outer electron
- 1 mark for linking same number of outer electrons to similar properties/reactions
- 1 mark for sodium and potassium electronic structures (can be stated or shown in diagrams)
Example 3: Drawing and interpreting electronic structure
Question: (a) Draw an electron shell diagram for a chlorine atom. [2 marks] (b) State which group of the periodic table chlorine is in and explain your answer. [2 marks]
Solution:
(a) Diagram showing:
- Nucleus labelled Cl or 17p, 18n ✓
- 2 electrons in first shell, 8 in second shell, 7 in third shell ✓
(b) Group 7 or Group 17 ✓
Chlorine has 7 electrons in its outer shell / the group number equals the number of outer electrons ✓
Mark scheme notes: For part (a), marks awarded for correct nucleus representation and correct electron distribution. For part (b), stating the group and correctly linking to electronic structure.
Common mistakes and how to avoid them
Mistake: Exceeding shell capacities (e.g., writing 2,10,8 for calcium instead of 2,8,8,2)
- Fix: Remember the maximum for shells 2 and 3 is 8 electrons at GCSE level. Check your answer adds up to the atomic number.
Mistake: Forgetting to adjust electron numbers for ions
- Fix: Positive ions have fewer electrons than protons (subtract electrons). Negative ions have more electrons than protons (add electrons). For example, Na⁺ has 11 protons but only 10 electrons.
Mistake: Confusing group numbers with outer electrons for transition metals
- Fix: The group number = outer electrons rule only applies to Groups 1, 2, and 13-18, not transition metals in the middle block.
Mistake: Starting to fill the second shell before the first shell is full
- Fix: Shells fill from innermost to outermost. Always fill lower energy shells first. First shell must have 2 electrons before any go in the second shell.
Mistake: Drawing electron shell diagrams with electrons inside shells rather than on the shell line
- Fix: Position electrons precisely on the circle representing each shell, spaced evenly around it.
Mistake: Writing electronic structure for elements beyond calcium without being careful
- Fix: At GCSE, you're only expected to know the pattern for elements 1-20. If asked about heavier elements, you may be given information or should state the pattern becomes more complex.
Exam technique for electronic structure questions
"State" or "Give" questions require brief answers without explanation. For electronic structure, simply write the numbers (e.g., "2,8,7"). Don't waste time explaining unless asked. (1 mark typically)
"Draw" questions require clear diagrams. Use a ruler for shells if possible, space electrons evenly, and label the nucleus. Check your electron count matches the atomic number. (Usually 2 marks: 1 for nucleus, 1 for electron distribution)
"Explain" questions need you to link electronic structure to properties or reactivity. Always mention the number of outer shell electrons and how this affects reactivity. Use connecting words like "therefore" or "because". (Typically 2-3 marks with 1 mark per valid point)
Check your arithmetic: Always verify that the sum of electrons in your electronic structure equals the atomic number (for atoms) or atomic number ± charge (for ions). This simple check catches most errors and takes only seconds.
Quick revision summary
Electronic structure describes how electrons are arranged in shells around the nucleus. Shells fill from innermost outwards with capacities of 2, 8, 8 (for elements 1-20 at GCSE). Write structures as comma-separated numbers (e.g., sodium is 2,8,1). The periodic table reflects electronic structure: group number indicates outer electrons, period number indicates number of shells. Elements react to achieve stable noble gas configurations with full outer shells. Same group elements have similar properties because they have the same number of outer electrons.