UNIT 2 CHEMISTRY · PEARSON AAQ · A1.2.1
Energy Levels, Subshells and Electronic Orbitals
FREE CHEMISTRY REVISION · TOPIC 4
Shell. Subshell. Orbital.
Build a precise picture of how electron states are organised, what orbital shapes mean and how to calculate their capacities.
New Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · A1.2.1
What you will learn
- Distinguish principal shells, subshells and individual atomic orbitals.
- Relate the principal quantum number n to the allowed subshells.
- Recall the orbital counts and maximum electron capacities of s, p and d subshells.
- Interpret scientifically accurate s and p orbital diagrams.
- Calculate capacities without confusing available states with actual electron occupancy.
Mapped to Unit 2 A1.2.1 in the current Pearson AAQ Applied Science specification, Issue 7 (May 2026), printed page 36. Detailed electron-filling rules are reserved for A1.2.2; atom/ion configuration and electron-in-box practice are reserved for A1.2.3.
1. Three levels of organisation
A simple shell model is useful, but electrons do not travel around the nucleus on fixed circular tracks. Quantum mechanics describes allowed electron states and the probability of finding an electron in different regions of space.
Principal shell: labelled by n
A shell groups states with the same principal quantum number, n = 1, 2, 3, … . It is also called a principal energy level. Higher-n states generally extend farther from the nucleus; this is not a set of hard, non-overlapping spherical layers.
Subshell: labelled by a number and letter
A subshell is a set of orbitals within a shell with the same subshell type: s, p, d or f. For example, 2p means the p subshell in shell n = 2. The 2 does not count its orbitals.
Orbital: one allowed spatial state
An atomic orbital is an allowed one-electron quantum state with a characteristic spatial probability distribution. At this level it is often described as a region around the nucleus where there is a high probability of finding an electron. A single orbital can accommodate at most two electrons, with opposite spins. It may also be empty or singly occupied.
A shell contains subshells.
A subshell contains one or more orbitals.
Each orbital holds a maximum of 2 electrons.
An orbital is not an electron: an unoccupied orbital is still an available state. Electron spin is a quantum property, not a literal tiny ball spinning on its axis. The two-electron capacity is all we need here; applying the filling rules is the next topic.
2. Which subshells belong to each shell?
The value of n determines which subshell types are allowed. Shell n has n subshell types, beginning with s, then p, d and f as n increases.
| Shell | Subshells | Orbitals in the shell | Maximum electrons |
|---|---|---|---|
| n = 1 | 1s | 1 | 2 |
| n = 2 | 2s, 2p | 1 + 3 = 4 | 8 |
| n = 3 | 3s, 3p, 3d | 1 + 3 + 5 = 9 | 18 |
| n = 4 | 4s, 4p, 4d, 4f | 1 + 3 + 5 + 7 = 16 | 32 |
There is no 1p, 2d or 3f subshell. The s type first occurs at n = 1; p at n = 2; d at n = 3. The f type first occurs at n = 4 and is included only to complete the shell pattern; detailed f-orbital treatment is outside this lesson.
Total orbitals in shell n = n².
Maximum electron capacity of shell n = 2n².
Capacity is not filling order. The third shell can accommodate 18 electrons, even though the simple arrangements of some early elements show only eight in that shell. This does not mean every 3d state fills before 4s is occupied. The order of occupancy is addressed in A1.2.2.
3. Count s, p and d orbitals
| Subshell type | Orbitals | Maximum electrons | Examples |
|---|---|---|---|
| s | 1 | 2 | 1s, 2s, 3s |
| p | 3 | 6 | 2p, 3p, 4p |
| d | 5 | 10 | 3d, 4d, 5d |
The letter determines the orbital count. A 2p subshell and a 3p subshell each contain three orbitals and each have a maximum capacity of six electrons. Their shell numbers, sizes and energies differ; their capacities do not.
A p subshell contains px, py and pz orbitals, oriented along three mutually perpendicular axes. A d subshell contains five orbitals. Four conventional d-orbital shapes have four main lobes; the dz² orbital has two main lobes and a ring. Detailed d shapes are not needed for the capacity calculation: five orbitals still give ten electrons.
Equal energy within a subshell: in an isolated atom without an external field, orbitals of the same subshell have equal energy (they are degenerate). Do not confuse three equal-energy 2p orbitals with three different subshells.
4. Understand s and p orbital shapes
An s orbital has spherical symmetry
The probability density for an s orbital is the same in every direction at a given distance from the nucleus. Its boundary surface is drawn as a sphere centred on the nucleus. A flat circle is only the two-dimensional representation of that sphere.
Higher s orbitals, such as 2s, retain spherical symmetry but have radial structure, including regions of zero probability density called radial nodes. A single shaded sphere does not show every detail of a higher s orbital; the diagram above deliberately uses 1s.
A p orbital has two lobes
A typical p-orbital boundary surface is described as dumbbell-shaped. The two lobes belong to one orbital, with the nucleus at their centre. A nodal plane through the nucleus separates the lobes. For 2px, that plane is perpendicular to the x axis.
Do not place “one electron in each lobe”. Electrons occupying the same orbital share its overall spatial probability distribution; the drawing does not assign either electron to a particular side. The node refers to this orbital’s probability density, not to the absence of the nucleus or of all other orbitals there.
5. Shells and subshell energies
Electron energies are quantised: only certain allowed values occur in an atom. A shell label alone does not specify every orbital energy in an atom containing several electrons. In such atoms, subshells within a shell generally have different energies; for example, 2s is lower in energy than 2p.
In the simple hydrogen-atom model, subshells with the same n have the same energy. The additional splitting between subshells in atoms containing several electrons reflects electron interactions. This is why it is useful to distinguish shells from subshells.
No universal shell-by-shell ladder: states from different principal shells can overlap in energy, and energies depend on the species and occupancy. Do not infer a full filling sequence just by arranging the leading numbers. We will apply the relevant filling rules in A1.2.2.
6. Worked examples
Example 1 · Decode 3p
- The number 3 identifies shell n = 3.
- The letter p identifies the subshell type.
- Every p subshell contains 3 orbitals.
- Maximum capacity = 3 × 2 = 6 electrons.
This information does not tell you how many electrons are actually present.
Example 2 · Capacity of the third shell
- List the allowed subshells: 3s, 3p and 3d.
- Count orbitals: 1 + 3 + 5 = 9.
- Multiply by two: 9 × 2 = 18 electrons maximum.
- Check using 2n²: 2 × 3² = 18.
Example 3 · Read a supplied label, 3d⁷
The 3d subshell is in shell n = 3 and contains five orbitals. The superscript 7 means seven electrons are present; it does not mean seven orbitals. Its capacity is ten, so ten − seven = three additional electrons would bring it to full capacity. How those seven electrons are distributed between the five orbitals is reserved for the filling-rule and configuration topics.
Example 4 · Check an impossible claim
A student says, “The 2d subshell contains five orbitals.” The orbital count would be correct for a d subshell, but 2d does not exist: shell n = 2 permits only 2s and 2p. Always check that a subshell is allowed before calculating its capacity.
Example 5 · One p orbital versus a p subshell
One 2p orbital has two lobes and holds at most 2 electrons. The complete 2p subshell has three such orbitals and holds at most 6 electrons. Count orbitals, not lobes.
7. Common misconceptions and exam tips
- “An orbital is an orbit.” It describes a quantum state and probability distribution, not a fixed path.
- “Every orbital holds six electrons.” Every single orbital holds at most two; a p subshell holds six across three orbitals.
- “Two p lobes mean two orbitals.” The two lobes belong to one orbital.
- “3p has three electrons.” The 3 is the shell label. A superscript, such as 3p⁴, gives the number of electrons present.
- “All shells hold eight electrons.” The maximum is 2n²: 2, 8, 18, 32 for the first four shells.
- “A capacity table tells me the filling order.” Capacity and occupancy are different. Use the rules in A1.2.2 to determine occupancy.
- “The orbital boundary is a wall.” The drawn surface encloses a chosen high probability; probability extends beyond it.
Exam tip: write the object you are counting: shell, subshell, orbital or electron. For a capacity question, show “number of orbitals × 2”. For a shape question, use spherical for s and two-lobed/dumbbell-shaped for a typical p orbital, and label the nucleus.
8. Check your understanding
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Your revision summary
- Shell: principal quantum number n.
- Subshell: number plus letter, such as 2p.
- Orbital: one spatial quantum state, maximum two electrons with opposite spins.
- s, p and d subshells have 1, 3 and 5 orbitals; capacities 2, 6 and 10.
- Shell n has n² orbitals and maximum capacity 2n².
- s orbitals have spherical symmetry; one p orbital has two lobes.
- Capacity does not specify the actual filling order or electron occupancy.
Without notes, explain why a 3p subshell can hold six electrons even though a single p orbital has only two lobes.