Electron structure

In a nutshell

At GCSE you drew electrons in shells as 2, 8, 8. At A level that picture is refined: each shell is split into sub-shells (s, p, d), and each sub-shell is made of orbitals that hold up to two electrons each.

This subtopic teaches you how electrons fill those orbitals, and how to write the electron configuration of any atom up to krypton (Z = 36) and of s- and p-block ions. Almost everything later in the course (bonding, ionisation energies, the periodic table) is explained with this model.

Assumed knowledge: Atomic structure and isotopes (atomic number, and how many electrons an atom or ion has).

Core content

Shells and how many electrons they hold

Electrons sit in shells (also called energy levels), numbered n = 1, 2, 3, 4 outwards from the nucleus. The higher the shell number, the further from the nucleus and the higher the energy.

The maximum number of electrons each of the first four shells can hold is:

The pattern is 2n² (2 × 1, 2 × 4, 2 × 9, 2 × 16), which is a handy way to check yourself.

Watch out: the third shell holds 18, not 8. The GCSE "2, 8, 8" pattern only describes how the first 18 electrons fill, not how many the third shell can hold.

Atomic orbitals and sub-shells

An atomic orbital is a region around the nucleus that can hold up to two electrons, with opposite spins.

Two electrons sharing one orbital must spin in opposite directions. In diagrams this is drawn as one up arrow and one down arrow: ↑↓.

Shapes of s- and p-orbitals

• An s-orbital is spherical, centred on the nucleus.
• A p-orbital is dumbbell-shaped: two lobes, one on each side of the nucleus.
• A p-sub-shell has three p-orbitals at right angles to each other, lying along the x, y and z axes (p_x, p_y, p_z).

Sub-shells

Orbitals of the same type in the same shell form a sub-shell. You need to know how many orbitals each sub-shell has and how many electrons it holds:

The electron count is always 2 × the number of orbitals, because each orbital holds two.

Which sub-shells are in each shell

The fourth shell also has an f-sub-shell (seven orbitals, 14 electrons). You only need it to explain where the 32 comes from.

Still don't get it? · shells, sub-shells and orbitals

Think of a block of flats.

• Each floor is a shell. Floor 1 is nearest the ground (the nucleus).
• Each floor is divided into flats of different types: a small s-flat, a bigger p-flat, an even bigger d-flat. These are the sub-shells.
• Each flat is divided into rooms. An s-flat has 1 room, a p-flat has 3, a d-flat has 5. These are the orbitals.
• Every room has exactly two beds, and the two people in a room always sleep head-to-toe (opposite spins).

Now count. A p-flat has 3 rooms, so it sleeps 3 × 2 = 6. A d-flat has 5 rooms, so it sleeps 10. Floor 2 has an s-flat and a p-flat, so it holds 2 + 6 = 8.

The usual mix-up is between a room and a flat: "a p-orbital holds 6 electrons" is wrong. One p-orbital holds 2; the p-sub-shell (all three p-orbitals) holds 6.

In exam wording: an orbital is a region around the nucleus that can hold up to two electrons, with opposite spins; a sub-shell is a set of orbitals of the same type (s, p or d) in one shell.

How the model developed

The GCSE "2, 8, 8" picture is the Bohr model, with electrons in circular shells. New evidence, such as the pattern in first ionisation energies across a period (OCR 3.1.1), could not be explained by shells alone. The model was refined to include sub-shells and orbitals: a good example of a scientific model being improved as evidence builds up.

Filling orbitals in order of increasing energy

Electrons fill orbitals in order of increasing energy: the lowest-energy sub-shell available fills first.

The order you need, up to Z = 36:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p

The key point: the 4s sub-shell is lower in energy than the 3d sub-shell, so 4s fills before 3d.

• After argon (18 electrons, 3p full), the next electron goes into 4s, not 3d. So potassium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
• Calcium fills 4s (4s²).
• Scandium to zinc then fill 3d.
• Gallium to krypton fill 4p.

Still don't get it? · why 4s fills before 3d

Imagine a car park where drivers always take the cheapest free space.

Lower floors are usually cheaper. But on this car park, one row of spaces on floor 4 (call it "4s") is cheaper than the expensive row on floor 3 (call it "3d"). So once the cheap rows on floors 1 to 3 are full, the next drivers park in 4s, and only then start using 3d.

Step by step, for electrons:

1. Electrons always go into the lowest-energy space available.
2. The energy of a sub-shell depends on its shell number and its type (s, p or d), so the order is not just "shell 1, then shell 2, then shell 3".
3. Up to 3p, the order happens to follow the shells.
4. Then 4s is lower in energy than 3d, so electrons 19 and 20 go into 4s.
5. Only after 4s is full does 3d start to fill (scandium onwards), followed by 4p.

In exam wording: orbitals fill in order of increasing energy, and the 4s sub-shell is lower in energy than the 3d sub-shell, so it fills first.

Same-energy orbitals: singly before pairing

The orbitals in one sub-shell (the three p-orbitals, or the five d-orbitals) all have the same energy.

When orbitals have the same energy, electrons occupy them singly before pairing.

• Electrons repel each other, so they spread out into separate empty orbitals first.
• Unpaired electrons in separate orbitals are drawn with the same spin (all arrows up).
• Only when every orbital in the sub-shell has one electron does the next electron pair up, with opposite spin.

Electrons-in-boxes for nitrogen (Z = 7), 1s² 2s² 2p³

The three 2p electrons go one into each 2p orbital, so nitrogen has three unpaired electrons.

Electrons-in-boxes for oxygen (Z = 8), 1s² 2s² 2p⁴

The fourth 2p electron has nowhere empty to go, so it pairs up. Oxygen has two unpaired electrons.

Rules for drawing electrons-in-boxes:

• One box = one orbital.
• One arrow = one electron.
• Two arrows in a box must point in opposite directions (opposite spins).
• Draw every orbital of a sub-shell, including empty ones, and fill each sub-shell singly before pairing.

Electron configurations of atoms

An electron configuration shows how many electrons are in each sub-shell, written in sub-shell notation.

For oxygen: 1s² 2s² 2p⁴

• The big number (2) is the shell.
• The letter (p) is the sub-shell.
• The superscript (4) is the number of electrons in that sub-shell.

Method for an atom with atomic number Z

1. Take Z from the Periodic Table on the Data Sheet: a neutral atom has Z electrons.
2. Fill sub-shells in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, putting in the maximum (s = 2, p = 6, d = 10) until you run out.
3. Check: the superscripts must add up to Z.

Examples

Notice that the configurations above are written in shell order (3d next to the other shell-3 sub-shells, before 4s), which is how the OCR specification writes them. The filling order is still 4s before 3d: it is only the written order that groups by shell.

Shorthand. You may see configurations shortened using the previous noble gas in square brackets, for example Ca as [Ar] 4s². It is a useful check, but if a question asks for the full electron configuration, write every sub-shell out from 1s.

Two exceptions: chromium and copper

Two atoms up to Z = 36 do not follow the filling order. Learn them as facts:

In each, one 4s electron moves into 3d, giving a half-filled (3d⁵) or completely filled (3d¹⁰) d-sub-shell, which is more stable.

Electron configurations of ions

For ions, start from the atom's configuration, then remove or add electrons to match the charge. At AS you only need s- and p-block ions up to Z = 36.

Positive ions: electrons are removed from the outermost shell first (the highest shell number), starting with its highest sub-shell.

• Na: 1s² 2s² 2p⁶ 3s¹ → Na⁺: 1s² 2s² 2p⁶
• Mg: 1s² 2s² 2p⁶ 3s² → Mg²⁺: 1s² 2s² 2p⁶
• Al: 1s² 2s² 2p⁶ 3s² 3p¹ → Al³⁺: 1s² 2s² 2p⁶ (3p electron and both 3s electrons lost)

Negative ions: electrons are added to the next available orbitals, following the normal filling order.

• O: 1s² 2s² 2p⁴ → O^2-: 1s² 2s² 2p⁶
• Cl: 1s² 2s² 2p⁶ 3s² 3p⁵ → Cl^-: 1s² 2s² 2p⁶ 3s² 3p⁶

Many s- and p-block ions end up with the same configuration as a noble gas. For example, K⁺, Ca²⁺, Cl^- and S^2- all have 18 electrons: 1s² 2s² 2p⁶ 3s² 3p⁶, the same as argon.

Quick check for any ion: number of electrons = Z - (ionic charge), so a 2+ ion has Z - 2 electrons and a 2- ion has Z + 2.

Worked examples

This subtopic has no numerical calculations. The skill is deduction: building a configuration from Z (and charge), then checking the electron total.

Worked example 1: phosphorus atom (Z = 15), with electrons-in-boxes

1. Phosphorus has 15 electrons.
2. Fill in order: 1s² (2), 2s² (4), 2p⁶ (10), 3s² (12), then 3 left for 3p.
3. Configuration: 1s² 2s² 2p⁶ 3s² 3p³
4. Check the total:

$2 + 2 + 6 + 2 + 3 = 15$

5. The three 3p electrons occupy the three 3p orbitals singly, so phosphorus has three unpaired electrons:

Where marks go: drawing the 3p as ↑↓ ↑ (one pair plus one single) loses the mark. Same-energy orbitals fill singly first.

Worked example 2: iron atom (Z = 26), filling 4s before 3d

1. Iron has 26 electrons.
2. Up to 3p⁶ uses 18 electrons (the argon core).
3. The next sub-shell is 4s (lower in energy than 3d): 4s² uses 2 more, making 20.
4. The remaining 6 go into 3d: 3d⁶.
5. Written in shell order: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
6. Check:

$2 + 2 + 6 + 2 + 6 + 6 + 2 = 26$

7. The 3d sub-shell has 5 orbitals and 6 electrons: one in each orbital, then the sixth pairs. So there is one pair and four unpaired 3d electrons.

Where marks go: the common wrong answer is ...3p⁶ 3d⁸, which fills 3d before 4s.

Worked example 3: identify the element from a configuration

An atom has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p³. Which element is it?

1. Add the superscripts to get the number of electrons, which equals Z for a neutral atom:

$2 + 2 + 6 + 2 + 6 + 10 + 2 + 3 = 33$

2. Z = 33 is arsenic, As.
3. Cross-check: it ends in 4p³, so it is in the p-block, in the same group as nitrogen and phosphorus (which end in p³).

Worked example 4: ions, S^2- (Z = 16) and Ga³⁺ (Z = 31)

Sulfide ion, S^2-:

1. S atom: 1s² 2s² 2p⁶ 3s² 3p⁴ (16 electrons).
2. A 2- charge means 2 electrons gained: 16 + 2 = 18 electrons.
3. They go into the half-empty 3p: S^2-: 1s² 2s² 2p⁶ 3s² 3p⁶

Gallium ion, Ga³⁺:

1. Ga atom: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹ (31 electrons).
2. A 3+ charge means 3 electrons lost: 31 - 3 = 28 electrons.
3. Remove from the outermost shell (shell 4): first the 4p electron, then both 4s electrons. The full 3d¹⁰ stays.
4. Ga³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰
5. Check:

$2 + 2 + 6 + 2 + 6 + 10 = 28$

Where marks go: taking electrons out of 3d (giving ...3d⁷ 4s² 4p¹) is wrong. Electrons are lost from the outermost shell.

Common exam mistakes

• Putting electron 19 into 3d. Potassium is ...3p⁶ 4s¹, not ...3p⁶ 3d¹. The 4s sub-shell is lower in energy than 3d, so it fills first.
• Pairing before every orbital has one electron. Nitrogen's 2p is ↑ ↑ ↑, not ↑↓ ↑ and an empty box. Same-energy orbitals are occupied singly before pairing.
• Two arrows pointing the same way in one box. Two electrons in one orbital must have opposite spins (↑↓).
• Mixing up orbital and sub-shell. "A p-orbital holds 6 electrons" is wrong: one p-orbital holds 2; the p-sub-shell (three orbitals) holds 6.
• Leaving "opposite spins" out of the orbital definition. The full idea is a region around the nucleus that holds up to two electrons with opposite spins.
• Calling an s-orbital a circle. Orbitals are three-dimensional: an s-orbital is spherical, a p-orbital is dumbbell-shaped.
• Saying the third shell holds 8 electrons. That is the GCSE habit. Shell 3 holds 18 (3s, 3p and 3d).
• Writing GCSE notation. "2, 8, 1" for sodium earns nothing when sub-shell notation is asked for: write 1s² 2s² 2p⁶ 3s¹.
• Getting chromium and copper wrong. They are 3d⁵ 4s¹ and 3d¹⁰ 4s¹, not 3d⁴ 4s² and 3d⁹ 4s².
• Adding or removing electrons the wrong way for ions. A positive ion has fewer electrons than the atom; a negative ion has more. Cl^- has 18 electrons, not 16.
• Using shorthand when the full configuration is asked for. [Ne] 3s¹ is not a full electron configuration; write it out from 1s.
• Not checking the total. The superscripts must add up to the number of electrons (Z for an atom, Z minus the charge for an ion). Adding them up catches most slips.

Key definitions

OCR's specification describes these terms rather than requiring formal definitions, so the wording below stays close to the specification's own.

• Atomic orbital: a region around the nucleus that can hold up to two electrons, with opposite spins.
• Shell (energy level): a main energy level, numbered n = 1, 2, 3, 4 outwards from the nucleus and made up of one or more sub-shells; the first four shells hold up to 2, 8, 18 and 32 electrons.
• Sub-shell: a set of orbitals of the same type (s, p or d) within a shell. An s-sub-shell has 1 orbital (2 electrons), a p-sub-shell 3 orbitals (6 electrons) and a d-sub-shell 5 orbitals (10 electrons).
• Filling order: orbitals fill in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p (4s is lower in energy than 3d).
• Orbitals of the same energy: electrons occupy them singly before pairing.
• Electron configuration: the arrangement of electrons in the sub-shells of an atom or ion, written in sub-shell notation, e.g. oxygen 1s² 2s² 2p⁴.

Specification

• I can state the maximum number of electrons that can fill each of the first four shells (2, 8, 18, 32). (2.2.1(a))
• I can describe an atomic orbital as a region around the nucleus that can hold up to two electrons, with opposite spins. (2.2.1(b)(i))
• I can describe (and sketch) the shapes of s- and p-orbitals. (2.2.1(b)(ii))
• I can state the number of orbitals in s-, p- and d-sub-shells, and the number of electrons each sub-shell can hold. (2.2.1(b)(iii))
• I can explain how the model of electron structure has been refined from shells to sub-shells and orbitals as evidence developed. (2.2.1(b), HSW1 and HSW7)
• I can fill orbitals for the first three shells and the 4s and 4p orbitals in order of increasing energy, knowing 4s fills before 3d. (2.2.1(c)(i))
• I can apply the rule that orbitals of the same energy are occupied singly before pairing, and draw electrons-in-boxes diagrams. (2.2.1(c)(ii))
• I can deduce the electron configuration of any atom up to Z = 36 in sub-shell notation, including chromium and copper. (2.2.1(d)(i))
• I can deduce the electron configuration of s- and p-block ions up to Z = 36 from the atomic number and ionic charge. (2.2.1(d)(ii))

Related notes

• Atomic structure and isotopes: counting the electrons in atoms and ions
• Bonding and structure: outer-shell electrons in ionic and covalent bonding, and the shapes of molecules
• Periodicity: s-, p- and d-blocks, and how ionisation energies give evidence for shells and sub-shells
• Transition elements (A level only): configurations of d-block ions