s, p, d, f Block Elements Revision Notes | A Level AQA Chemistry

1. Why the Periodic Table Has Four Distinct Regions

The Periodic Table is not arranged randomly. Every element sits in a precise position determined by its proton number — and that position tells you exactly which type of subshell its outermost electrons occupy. This is the foundation of s, p, d, and f block classification.

Think of the Periodic Table as a map of electron configurations. As you move from left to right across a period, you are filling subshells in a fixed order: s, then d (in the middle), then p on the right. The f block sits in its own region, usually shown below the main table. Each of these regions is a block.

The AQA specification states this directly: an element is classified as s, p, d, or f block according to its position in the Periodic Table, which is determined by its proton number. That means the classification is not about the element's most stable ion or its common oxidation state — it is about where the element sits and which subshell is being filled by its electrons in the ground state.

2. The s Block: Groups 1 and 2

The s block occupies the two leftmost columns of the Periodic Table: Group 1 (the alkali metals) and Group 2 (the alkaline earth metals). In these elements, the highest-energy electrons are filling an s subshell — specifically the ns subshell, where n is the period number.

Notice that helium is the classic exception to watch for. Its chemical behaviour resembles the noble gases, so it is placed in Group 18 on most Periodic Tables. However, its electron configuration is 1s² — there is no p subshell involved — so it is correctly classified as an s block element. AQA examiners have tested this distinction before.

3. The p Block: Groups 13 to 18

Cross to the right-hand side of the Periodic Table and you reach the p block, spanning Groups 13 to 18. Here, the highest-energy electrons are filling a p subshell. A p subshell holds a maximum of 6 electrons, which is why the p block is six groups wide.

The p block contains a huge variety of elements: metals such as aluminium and lead, non-metals such as nitrogen, oxygen, and chlorine, metalloids such as silicon, and the noble gases (except helium). What they share is that their outermost electrons are in an np subshell.

The pattern is clear: as you move from Group 13 to Group 18, you are adding one electron at a time to the np subshell. The proton number increases by one with each step, and the position in the table shifts one group to the right. This is exactly what AQA means when it says position is determined by proton number.

4. The d Block: The Transition Metals (Groups 3 to 12)

Between the s block and the p block sits the d block — the wide central section of the Periodic Table, spanning Groups 3 to 12. These are the d block elements, which include the transition metals. Their characteristic feature is that electrons are being added to a d subshell. A d subshell holds a maximum of 10 electrons, which is why the d block is ten groups wide.

There is a subtlety here worth examining carefully. The d subshell being filled in Period 4 is the 3d subshell, even though these elements are in Period 4. According to the Aufbau principle, the 4s subshell fills before the 3d subshell, as it has lower energy in the neutral atom before the 3d begins filling. So for Period 4 d block elements, the configuration follows the pattern [Ar] 3d$^{x}$ 4s$^{2}$ (or 4s$^{1}$ for chromium and copper, which have well-known exceptions).

The copper and chromium exceptions arise because a half-filled (3d⁵) or fully filled (3d¹⁰) d subshell is particularly stable. One electron is promoted from the 4s subshell to achieve this configuration. You are expected to know these two exceptions at A Level.

5. The f Block: Lanthanides and Actinides

The f block is the region most students encounter least often, but it follows the same logic. These are the lanthanides (Period 6, proton numbers 57–71) and the actinides (Period 7, proton numbers 89–103). Their highest-energy electrons are filling an f subshell, which holds a maximum of 14 electrons — hence the f block is 14 elements wide.

On most printed Periodic Tables, the f block is detached and shown below the main body of the table. This is a practical layout choice, not a chemical one. If the f block were inserted in its correct position, it would sit between Groups 2 and 3, making the table extremely wide. The proton numbers still run consecutively — the layout is the only thing that changes.

For AQA A Level, you are not required to memorise f block electron configurations in detail, but you must know that f block elements exist, where they sit relative to the rest of the table, and that their classification follows the same principle: the subshell being filled determines the block.

6. Mapping the Blocks: A Visual Overview

Bringing all four blocks together, the Periodic Table can be divided into four distinct regions. The key is to read the table from left to right across each period and track which subshell is being filled at each point.

Notice the order in which blocks appear across a period: s block first, then d block (from Period 4 onwards), then p block. The f block inserts between s and d from Period 6 onwards. This order reflects the order in which subshells fill according to the Aufbau principle.

7. Proton Number as the Determining Factor

The AQA specification is explicit: block classification is determined by proton number. This is worth unpacking. As proton number increases by one, one more electron is added to the atom (in the ground state). The subshell that receives that electron determines which block the element belongs to.

Consider the transition from calcium (Z = 20, s block) to scandium (Z = 21, d block). Calcium has the configuration [Ar] 4s² — the 4s subshell is full. The next electron added (for scandium) goes into the 3d subshell, giving [Ar] 3d¹ 4s². That single extra proton shifts the element from s block to d block. The position in the Periodic Table changes, and so does the block.

Worked example: Classifying an Element by Block from its Proton Number

Classify germanium (Ge, Z = 32) as s, p, d, or f block. Justify your answer using its electron configuration.

1. Z = 32: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p² — Fill subshells in order using the Aufbau principle. After the 3d subshell is complete at zinc (Z = 30), the 4p subshell begins filling from gallium (Z = 31) onwards.
2. The last electrons added are in the 4p subshell:...4s² 4p² — Moving from Z = 31 (Ga, 4p¹) to Z = 32 (Ge, 4p²), the additional electron enters the 4p subshell. This is the subshell being filled at this point in the Periodic Table.
3. Highest-energy subshell = 4p → p block — Because the subshell being filled is a p subshell, germanium is classified as a p block element.
4. Ge is in Group 14, Period 4 — Group 14 falls within Groups 13–18, which is the p block region. This confirms the classification. Proton number 32 places germanium in Period 4, Group 14 — unambiguously p block.

Answer: Germanium is a p block element. Its electron configuration is [Ar] 3d¹⁰ 4s² 4p², and the highest-energy electrons are in the 4p subshell.

8. Exam Technique and Common Mistakes

Block classification questions at A Level tend to be short — often just 1 or 2 marks — but they are frequently dropped because students rely on vague memory rather than applying the electron configuration logic. The following guidance will help you secure every available mark.

Block classification is one of those topics where a clear, methodical approach — write the configuration, identify the subshell, name the block — will always get you the marks. The logic is consistent across every element in the Periodic Table, from hydrogen (Z = 1) to the heaviest actinides. Proton number is the anchor; electron configuration is the tool; block classification is the conclusion.