The AQA GCSE periodic table explained
The periodic table is the chemist's map of every element, arranged in order of increasing atomic number. Elements in the same vertical column (a group) have similar chemical properties because they have the same number of electrons in their outer shell. Elements in the same horizontal row (a period) have the same number of electron shells. For AQA GCSE Chemistry, the periodic table sits at the heart of Paper 1 and underpins almost every topic on bonding, reactions, and structure.
This guide explains how Mendeleev built the modern table, what groups and periods tell you, the key trends in Groups 1, 7, and 0, and the exam-ready facts the AQA mark scheme wants to see.
Ordered by atomic number
Each element sits in order of its proton number, not its mass. Hydrogen is element 1, helium is element 2, and so on.
Groups share properties
Elements in the same group have the same number of outer electrons, which is why they react in similar ways.
Trends down a group
Reactivity increases down Group 1, decreases down Group 7. Outer shells get further from the nucleus.
How the periodic table is arranged
Elements are arranged in order of increasing atomic number (the number of protons in the nucleus). The table has 7 rows called periods and 18 columns, of which the main groups for GCSE are Groups 1, 2, 3 to 6 (the middle p-block), 7, and 0.
The group number tells you the number of electrons in the outer shell. For example, lithium is in Group 1 and has 1 outer electron. Chlorine is in Group 7 and has 7 outer electrons. The period number tells you how many electron shells the atom has.
Group 0 vs Group 8 The noble gases sit in the far right column. In AQA GCSE Chemistry this is called Group 0 (because they have 0 reactive outer electrons in terms of bonding behaviour). You may also see it labelled Group 8 in older textbooks. AQA mark schemes use Group 0.
How Mendeleev built the modern table
Dmitri Mendeleev published his first periodic table in 1869, arranging the 60 or so known elements in order of atomic mass. His genius was leaving gaps for elements he predicted but had not been discovered yet, and swapping the order of a few pairs to keep elements with similar properties in the same column.
Later, when the structure of the atom was understood, the table was reordered by atomic number rather than atomic mass. This fixed the awkward pairs Mendeleev had swapped (such as tellurium and iodine) and explained why his predictions worked. The modern table is in atomic number order, but it is still recognisably his.
| Group | Common name | Outer electrons | Reactivity trend |
|---|---|---|---|
| Group 1 | Alkali metals | 1 | Reactivity increases down the group |
| Group 2 | Alkaline earth metals | 2 | Reactivity increases down the group |
| Group 7 | Halogens | 7 | Reactivity decreases down the group |
| Group 0 | Noble gases | 8 (2 for He) | Unreactive, full outer shell |
Group 1 – the alkali metals
Group 1 contains lithium, sodium, potassium, rubidium, caesium, and francium. They all have 1 outer electron, which they lose easily to form a +1 ion. This makes them very reactive metals, especially with water and oxygen.
Reactivity increases as you go down the group. Lithium fizzes gently in water, sodium melts into a ball and zips around the surface, potassium ignites with a lilac flame. The reason is that the outer electron is further from the nucleus in larger atoms, so it is held less tightly and is easier to lose.
Why reactivity increases down Group 1 The outer electron is further from the positive nucleus in larger atoms. It is also shielded by more inner electron shells. Both effects reduce the attraction holding the outer electron, so it is lost more easily, and the metal reacts faster.
Group 7 – the halogens
Group 7 contains fluorine, chlorine, bromine, iodine, and astatine. They all have 7 outer electrons and gain 1 more to form a 1- ion when they react with metals. They are non-metals and exist as diatomic molecules (F2, Cl2, Br2, I2).
Reactivity decreases as you go down the group, the opposite of Group 1. The reason is the same idea in reverse: A larger atom has its outer shell further from the nucleus, so it is harder to attract an extra electron in. Fluorine is the most reactive halogen, iodine is the least reactive of the common ones.
| Halogen | State at room temperature | Colour |
|---|---|---|
| Fluorine (F2) | Gas | Pale yellow |
| Chlorine (Cl2) | Gas | Green-yellow |
| Bromine (Br2) | Liquid | Red-brown |
| Iodine (I2) | Solid | Dark grey, purple vapour |
Group 0 – the noble gases
Group 0 contains helium, neon, argon, krypton, xenon, and radon. They all have a full outer shell of electrons (8 for the heavier ones, 2 for helium), which makes them stable and unreactive. They exist as single atoms (monatomic) and do not form molecules under normal conditions.
Because they are unreactive, the noble gases are used where chemical inertness matters: Argon fills incandescent light bulbs to stop the filament burning, helium fills weather balloons because it is light and non-flammable, neon glows orange-red in advertising signs.
Trends across a period
Across a period from left to right, you go from metals on the left to non-metals on the right. The number of outer electrons increases by 1 each time. The atoms get slightly smaller because the nucleus has more protons, pulling the electrons in tighter.
A period 3 example: Sodium (Group 1, metal) is on the far left, then magnesium, aluminium, silicon, phosphorus, sulfur, chlorine (Group 7, reactive non-metal), and argon (Group 0, unreactive noble gas) on the far right.
Metals on the left, non-metals on the right The staircase line that runs from boron down to astatine roughly separates metals from non-metals. Elements right next to the line (like silicon and germanium) have some properties of both and are called metalloids.
Where students lose marks on periodic table questions
AQA examiner reports flag the same mistakes every series. The most common ones are writing the wrong direction for a reactivity trend (saying Group 7 reactivity increases down the group, when it decreases), and confusing the group number with the period number.
Another classic error is saying "the atom gets bigger so it is more reactive" without explaining why. The full mark scheme answer for Group 1 is: The outer electron is further from the nucleus and shielded by more inner shells, so it is held less tightly and is lost more easily.
Key facts to memorise for the exam
- Elements are arranged in order of increasing atomic number (proton number)
- Group number = number of outer electrons (except Group 0)
- Period number = number of electron shells
- Group 1 alkali metals: Reactivity increases down the group
- Group 7 halogens: Reactivity decreases down the group
- Group 0 noble gases: Full outer shell, unreactive, monatomic
- Mendeleev left gaps for undiscovered elements and predicted their properties
- Modern table is ordered by atomic number, not atomic mass
