Overview
What you’ll learn
- Describe the arrangement of the Periodic Table in periods and groups by increasing proton number.
- Relate group number to outer-shell electrons and ion charge, and explain why elements in a group have similar reactions.
- Describe the trends in Group I and Group VII and use them to predict the properties of unseen members.
- Describe halogen displacement reactions, including observations and explanations.
- Describe the characteristic properties of transition elements and the unreactivity of the noble gases.
Syllabus checklist (0620)
Statements are paraphrased and grouped by tier. Core is studied by everyone; Supplement is examined on the Extended papers only.
Core — every candidate
- 8.1Describe the Periodic Table as an arrangement of elements in periods and groups, in order of increasing proton number.
- 8.1Describe the change from metallic to non-metallic character across a period.
- 8.1Describe the relationship between group number and the charge of the ions formed by elements in that group.
- 8.1Explain similarities in the chemical properties of elements in the same group in terms of their electronic configurations.
- 8.1Explain how the position of an element in the Periodic Table can be used to predict its properties.
- 8.2Describe lithium, sodium and potassium as relatively soft metals, with trends down the group of decreasing melting point, increasing density and increasing reactivity.
- 8.2Predict the properties of other Group I elements, given information about the group.
- 8.3Describe chlorine, bromine and iodine as diatomic non-metals, with trends down the group of increasing density and decreasing reactivity.
- 8.3State the appearance of the halogens at room temperature and pressure: chlorine a pale yellow-green gas, bromine a red-brown liquid, iodine a grey-black solid.
- 8.3Describe and explain the displacement reactions of halogens with other halide ions.
- 8.3Predict the properties of other Group VII elements, given information about the group.
- 8.4Describe transition elements as metals with high densities, high melting points, coloured compounds and which often act as catalysts.
- 8.5Describe the Group VIII noble gases as unreactive, monatomic gases, and explain this unreactivity in terms of their electronic configurations.
Supplement — Extended candidates only
- 8.1Identify trends in groups, given information about the elements.
- 8.4Describe transition elements as having ions with variable oxidation numbers, including iron(II) and iron(III).
In 30 seconds
- 1The modern Periodic Table lists elements in order of increasing proton number; rows are periods, columns are groups.
- 2Elements in the same group share the same number of outer-shell electrons, so they have similar chemical properties.
- 3Across a period, elements change from metallic to non-metallic character.
- 4Group I metals get more reactive down the group; Group VII halogens get less reactive down the group.
- 5A more reactive halogen displaces the ion of a less reactive one from solution.
- 6Transition elements have high densities and melting points, coloured compounds, variable oxidation numbers and are often catalysts; noble gases are unreactive because their outer shells are full.
Notes · 1 of 4
How the Periodic Table is arranged
Order by proton number; columns share outer-shell electrons and therefore chemistry.
The Periodic Table arranges all the elements in order of increasing proton number (atomic number). The horizontal rows are periods and the vertical columns are groups. For the first twenty elements, the period number equals the number of occupied electron shells, and the group number (Groups I–VII) equals the number of electrons in the outer shell.
| Group | Outer-shell electrons | Ion formed | Example |
|---|---|---|---|
| I | 1 | 1+ (electron lost) | Na⁺ |
| II | 2 | 2+ | Mg²⁺ |
| III | 3 | 3+ | Al³⁺ |
| VI | 6 | 2− (two electrons gained) | O²⁻ |
| VII | 7 | 1− | Cl⁻ |
Moving across a period, the elements change from metallic on the left through metalloids to non-metallic on the right, as oxides change from basic to acidic. Elements in the same group have the same number of outer-shell electrons, which is why they show similar chemical properties — and why an element’s position lets you predict how it will behave before you ever test it.
Notes · 2 of 4
Group I — the alkali metals
Soft, low-density metals whose reactivity rises steeply down the group.
Group I — lithium, sodium, potassium and the metals below them — are the alkali metals. They are relatively soft (cuttable with a knife) and have low densities compared with most metals. Each has one outer-shell electron, which it loses easily, so they all form 1+ ions and react vigorously with water to produce an alkali and hydrogen.
| Property | Trend |
|---|---|
| Melting point | Decreases |
| Density | Increases |
| Reactivity with water | Increases |
- Lithium fizzes steadily as it moves on the water surface.
- Sodium melts into a ball and fizzes more vigorously.
- Potassium reacts so exothermically that the hydrogen produced ignites with a lilac flame.
These trends continue down the group: rubidium and caesium are denser, melt at lower temperatures and react even more violently than potassium. Exam questions often give data for an unseen member and expect you to extend the trend.
Notes · 3 of 4
Group VII — the halogens
Diatomic non-metals that get denser but less reactive down the group.
| Halogen | Appearance at r.t.p. |
|---|---|
| Chlorine, Cl₂ | Pale yellow-green gas |
| Bromine, Br₂ | Red-brown liquid |
| Iodine, I₂ | Grey-black solid |
All three are diatomic molecules with the formula X₂. Going down the group the melting and boiling points and densities increase, but the reactivity decreases — fluorine is the most reactive halogen and iodine the least of the four.
A more reactive halogen displaces a less reactive halogen from its aqueous halide ions. Chlorine displaces bromine from potassium bromide solution and iodine from potassium iodide solution; bromine displaces iodine; iodine displaces none of the others. The displaced halogen shows as a colour change in the solution.
- 1Chlorine into potassium bromide
The orange-brown colour of bromine appears: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂.
- 2Bromine into potassium iodide
The solution turns brown as iodine forms: Br₂ + 2I⁻ → 2Br⁻ + I₂.
- 3Iodine into potassium chloride
No reaction: iodine is less reactive than chlorine, so it cannot displace the chloride ions.
Notes · 4 of 4
Transition elements and noble gases
Dense, useful metals with coloured compounds — and the unreactive gases that end each period.
The transition elements occupy the block between Groups II and III. They are metals with high densities and high melting points. Their compounds are usually coloured — think of blue copper(II) sulfate, green iron(II) salts and purple potassium manganate(VII) — and they often act as catalysts, for example iron in the Haber process.
Supplement: unlike Group I metals, transition elements form ions with variable oxidation numbers. Iron forms iron(II), Fe²⁺, and iron(III), Fe³⁺, which is why the oxidation number must be included in the name of the compound.
| Property | Group I | Transition elements |
|---|---|---|
| Density | Low | High |
| Melting point | Decreases down the group | Generally high |
| Compounds | White or colourless | Often coloured |
| Catalytic behaviour | None | Many act as catalysts |
| Ion charge | Always 1+ | Variable |
The noble gases (Group VIII — helium, neon, argon, krypton, xenon, radon) are monatomic, colourless gases. They are strikingly unreactive because every atom already has a full outer shell of electrons, so they have no tendency to gain, lose or share electrons. This unreactivity makes them useful: argon provides an inert atmosphere in filament lamps and during welding.
Revise
Key terms
The words examiners expect you to use precisely.
- Period
- A horizontal row of the Periodic Table; the period number equals the number of occupied electron shells.
- Group
- A vertical column of the Periodic Table; group number (I–VII) equals the number of outer-shell electrons.
- Alkali metal
- A Group I metal: soft, low-density, forms 1+ ions and becomes more reactive down the group.
- Halogen
- A Group VII non-metal that exists as diatomic molecules, X₂, and becomes less reactive down the group.
- Displacement reaction
- A reaction in which a more reactive halogen removes a less reactive halogen from its halide ions.
- Diatomic
- Describes a molecule made of exactly two atoms, such as Cl₂, Br₂ and I₂.
Show 4 more key termsShow fewer key terms
- Transition element
- A metal in the block between Groups II and III: dense, high-melting, forms coloured compounds and often catalyses reactions.
- Noble gas
- A Group VIII element: a monatomic, unreactive gas with a full outer shell of electrons.
- Variable oxidation numbers
- The ability of transition elements to form ions of more than one charge, such as Fe²⁺ and Fe³⁺ (Supplement).
- Electronic configuration
- The arrangement of electrons in shells, e.g. 2,8,1 for sodium; it decides an element’s group and period.
Revise
Common mistakes
Where students most often lose marks in this topic — and what to do instead.
Mistake: Reversing the Group I reactivity trend.
Students often assume heavier metals are less reactive, as with halogens.
Instead: Group I reactivity **increases** down the group; Group VII reactivity decreases.
Mistake: Saying iodine displaces bromine from bromide solution.
Displacement only happens when the added halogen is the more reactive one.
Instead: Iodine is less reactive than bromine, so no displacement occurs.
Show 8 more mistakesShow fewer mistakes
Mistake: Confusing period number with group number.
Periods are rows and give the number of shells; groups are columns and give outer-shell electrons.
Instead: Sodium (2,8,1) is in period 3 and Group I.
Mistake: Listing low density and low melting points as transition-element properties.
Those belong to Group I; transition elements are dense and high-melting.
Instead: Transition elements: high density, high melting point, coloured compounds, catalytic behaviour.
Mistake: Explaining noble-gas unreactivity by saying "they are monatomic".
Being monatomic is a consequence of having a full outer shell, not the cause of unreactivity.
Instead: The full outer shell means there is no tendency to gain, lose or share electrons.
Mistake: Writing the halogen displacement equation without balancing charges.
Each halogen molecule needs two halide ions to supply both atoms.
Instead: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂ — two electrons and two ions per side.
Mistake: Assuming fluorine is the least reactive halogen because it is the lightest.
Reactivity decreases going down Group VII, so the lightest halogen is actually the most reactive.
Instead: Reactivity order: fluorine > chlorine > bromine > iodine.
Mistake: Trying to read a transition metal’s ion charge from a group number.
Transition elements sit in the block between Groups II and III and do not follow the outer-electron rule that works for Groups I–VII.
Instead: Transition ions come in several charges, which is why Roman numerals are written in their names.
Mistake: Describing the iodine formed in a displacement reaction as purple.
Purple describes iodine vapour or iodine in an organic solvent, not the colour that appears in water.
Instead: Displaced iodine makes the aqueous solution brown; displaced bromine makes it orange-brown.
Mistake: Writing halogens as single atoms (Cl, Br, I) in equations.
Halogens exist as diatomic molecules, so unbalanced atom counts make equations wrong before you start.
Instead: Write Cl₂, Br₂, I₂ and F₂ wherever the free halogen appears.
Check yourself
Quick quiz
18 questions with instant explanations. About 5 minutes.
Question 1 of 3
- Question 1: current question
- Question 2: not answered
- Question 3: not answered
State the ordering principle
In the modern Periodic Table, elements are arranged in order of increasing
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Practise
Topical questions
10 exam-style questions with model answers and mark schemes.
A set of structured and calculation questions matched to this topic, in three tiers, each with a model answer, mark scheme and examiner tip.
Practise topical questionsExam practice
Practice paper
A 30-mark paper set out like a Cambridge paper, with a separate mark scheme. Give yourself 40 min.
Question paper
Topic 8 Practice Paper
30 marks · 40 min · 6 pages
10 multiple-choice questions and structured questions with answer lines. Topic 8: The Periodic Table.
Mark scheme
PROTopic 8 Practice Paper — mark scheme
Answers · Marking points · 3 pages
Model answers with a mark for each creditworthy point — for marking at home or in class.
Which paper covers this topic?
Paper 1 / 2
Multiple choice (45 min, 40 marks)
Core and Extended versions
Paper 3 / 4
Theory, short and structured (1 h 15, 80 marks)
Core and Extended versions
Paper 5 / 6
Practical test or Alternative to Practical (40 marks)
Experimental skills
Want a full paper? Try the Paper 2-style multiple choice, Paper 4-style structured or Paper 6-style practical sets.
Questions
Frequently asked questions
How do I find the group and period of an element from its electronic configuration?
The number of occupied shells is the period, and the number of electrons in the outer shell is the group (for Groups I–VII). Sodium, 2,8,1, has three shells and one outer electron: period 3, Group I.
Why does reactivity increase down Group I but decrease down Group VII?
Group I metals react by losing their single outer electron, which gets easier the further the shell is from the nucleus. Halogens react by gaining an electron, and the growing atom attracts an extra electron less strongly, so reactivity falls down the group.
What are the halogen colours I must learn?
Chlorine is a pale yellow-green gas, bromine a red-brown liquid and iodine a grey-black solid at room temperature. In displacement reactions the displaced halogen shows as an orange-brown (bromine) or brown (iodine) colour in solution.
What makes transition elements different from Group I metals?
Transition elements are denser and higher-melting, form coloured compounds, often act as catalysts and have variable oxidation numbers such as Fe²⁺ and Fe³⁺. Group I metals are soft, low-density, form white or colourless compounds and always form 1+ ions.
Why are noble gases used in lamps and welding?
Noble gases have full outer electron shells, so they are extremely unreactive. Argon around a hot filament or a welding seam provides an atmosphere in which nothing can oxidise or burn.
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Topic 9: Metals