8. The Periodic Table
- Syllabus
- 0620–2026–2027
- Section
- 8
- Level
- —
Exam analysis
8.1 Arrangement of elements
- Syllabus
- 0620–2026–2027
- Topic
- 8.1
- Level
- —
Read the arrangement of the Periodic Table
The Periodic Table arranges elements in order of increasing proton number, also called atomic number. Each element therefore has one fixed position.
| Feature | Direction | What it tells you |
|---|---|---|
| period | horizontal row | atoms have the same number of occupied electron shells |
| group | vertical column | main-group atoms have the same number of outer-shell electrons |
| atomic number | increases from one element to the next | equals the number of protons in every atom of that element |
To identify a position, give both coordinates: for example, selenium is in Period 4 and Group VI. To compare two elements, first decide whether they share a row (period) or a column (group).
The table is ordered by proton number, not by relative atomic mass. Period number and group number describe different electron features and must not be interchanged.
Track metallic character across a period
Across a period from left to right, metallic character decreases and non-metallic character increases.
| Across a period | Left-hand elements | Right-hand elements |
|---|---|---|
| character | more metallic | more non-metallic |
| electron tendency | more likely to lose electrons and form positive ions | more likely to gain or share electrons |
| typical oxide character | more basic | more acidic |
If two elements are in the same period, the one farther left is normally more metallic. Thus sodium is more metallic than silicon, and silicon is more metallic than chlorine.
Do not use this left-to-right rule to compare elements in different periods without additional information. It describes a trend across one horizontal row.
Predict main-group ion charges
Main-group atoms form ions by losing or gaining outer-shell electrons to reach a full outer shell. The group number gives the number of outer electrons and therefore predicts the ion charge.
| Group | Outer electrons | Usual electron change | Typical ion charge |
|---|---|---|---|
| I | 1 | loses 1 | +1 |
| II | 2 | loses 2 | +2 |
| III | 3 | loses 3 | +3 |
| V | 5 | gains 3 | −3 |
| VI | 6 | gains 2 | −2 |
| VII | 7 | gains 1 | −1 |
Rubidium is in Group I, so it forms Rb⁺. Magnesium is in Group II, so it forms Mg²⁺. Sulfur is in Group VI, so it forms S²⁻.
The sign records electron movement: losing negative electrons gives a positive ion; gaining electrons gives a negative ion. Group IV elements do not normally form simple ±4 ions, and Group VIII noble gases do not usually form ions.
Explain similarities within a group
Elements in the same group have the same number of electrons in their outer shell. Because chemical reactions involve outer electrons, they have similar chemical properties.
| Same down a group | Changes down a group |
|---|---|
| number of outer-shell electrons | number of occupied electron shells |
| typical ion charge | atomic/proton number |
| broad reaction pattern and compound formula pattern | size, density, melting/boiling point and often reactivity |
Oxygen and sulfur are both in Group VI: each has six outer-shell electrons and commonly forms 2− ions. Calcium and strontium are both in Group II: each has two outer-shell electrons and commonly forms 2+ ions.
Similar chemical properties do not mean identical physical properties. Elements lower in a group have more electron shells, so their sizes and many measured properties change.
Predict properties from Periodic Table position
| Position clue | Prediction |
|---|---|
| period number | number of occupied electron shells |
| main-group number | number of outer-shell electrons |
| same group as a known element | similar reactions and typical compound formulae |
| left side | metallic behaviour, positive ions and often basic oxides |
| right side (except noble gases) | non-metallic behaviour and often acidic oxides |
| Group VIII | full outer shell and very low reactivity |
Locate the element, extract period and group information, then combine only justified clues. For an unknown below magnesium in Group II, predict a metal with two outer electrons, a 2+ ion, broadly similar reactions and a basic oxide.
Use predicted ion charge or valency to construct formulae. A Group II metal M and a Group VII non-metal X form MX₂; a Group I metal and a Group VI non-metal form M₂X.
Position supports broad predictions, not an exact melting point, colour or reaction rate unless a group trend or numerical data is supplied. State the periodic clue behind each prediction.
Identify a trend from group data
A group trend is a systematic change in a property as atomic number and position down the group increase.
| Step | Question to ask |
|---|---|
| order | Are the elements arranged from top to bottom of the group? |
| compare | Does the value increase, decrease or remain approximately constant? |
| check | Does the pattern hold across all supplied entries, allowing for stated anomalies? |
| predict | Should an unknown value lie above, below or between the known values? |
| justify | Which two or more data points establish the direction? |
If density rises 0.18 → 0.90 → 1.78 → 3.71 down a noble-gas group, identify increasing density. If melting points fall down a supplied group, predict a lower—not an invented exact—value for the next element unless interpolation is requested.
Do not import a memorised trend from another group. The objective asks you to identify the pattern in the information given; different properties and groups can change in different directions.
8.2 Group I properties
- Syllabus
- 0620–2026–2027
- Topic
- 8.2
- Level
- —
Describe the Group I trends
Lithium, sodium and potassium are Group I alkali metals. They are relatively soft metals and can be cut more easily than typical transition metals.
| Going down Group I | Trend | Lithium → sodium → potassium |
|---|---|---|
| melting point | decreases | highest → intermediate → lowest |
| density | generally increases | lowest → higher → generally higher |
| reactivity | increases | least reactive → more reactive → most reactive |
The increasing reactivity is seen in increasingly vigorous reactions. With water, each metal forms an alkaline metal hydroxide solution and hydrogen; potassium reacts more vigorously than sodium, which reacts more vigorously than lithium.
Keep the directions separate: down the group, melting point decreases, but density and reactivity increase. These are general trends; do not reverse density simply because melting point falls.
Predict properties of another Group I element
An unfamiliar Group I element should continue the pattern shown by the known elements, unless the supplied data reveal an exception.
| Step | Action | Example for rubidium below potassium |
|---|---|---|
| locate | place the unknown relative to known Group I elements | rubidium is below potassium |
| set direction | read the trend for each requested property | melting point ↓; density ↑; reactivity ↑ |
| bound | compare rather than invent unsupported precision | lower melting point and generally higher density than potassium |
| predict behaviour | extend the qualitative pattern | reacts more vigorously than potassium |
| check | apply every supplied datum and unit | a numerical estimate must fit the surrounding values |
If an unknown lies between two known elements, predict a value between their values. If it lies below the last known element, extrapolate in the established direction and use an approximate value only when the question asks for one.
Do not assume that every physical property rises or falls together. Make a separate prediction for melting point, density and reactivity, and justify each one from the given trend.
8.3 Group VII properties
- Syllabus
- 0620–2026–2027
- Topic
- 8.3
- Level
- —
Describe the Group VII trends
Chlorine, bromine and iodine are Group VII halogens. They are non-metals that exist as diatomic molecules: Cl₂, Br₂ and I₂.
| Going down Group VII | Trend | Chlorine → bromine → iodine |
|---|---|---|
| density | increases | lowest → intermediate → highest |
| reactivity | decreases | most reactive → less reactive → least reactive |
A more reactive halogen gains electrons more readily. Therefore chlorine reacts in situations where bromine or iodine may not, while iodine is less able to displace other halogens from their halides.
Group VII shows the opposite reactivity direction to Group I: halogen reactivity decreases down the group. Density, however, increases down Group VII.
State halogen appearances at r.t.p.
| Halogen | Colour at r.t.p. | State at r.t.p. | Complete appearance |
|---|---|---|---|
| chlorine | pale yellow-green | gas | pale yellow-green gas |
| bromine | red-brown | liquid | red-brown liquid |
| iodine | grey-black | solid | grey-black solid |
From chlorine to bromine to iodine, the colour becomes darker and the state changes from gas to liquid to solid.
When asked for appearance, give both colour and physical state, using room temperature and pressure as the condition.
Do not describe iodine crystals as purple at r.t.p.: solid iodine is grey-black. Purple describes iodine vapour, not the required room-condition solid.
Predict halogen displacement reactions
A more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide ions. Reactivity decreases F₂ > Cl₂ > Br₂ > I₂.
| Added halogen | chloride, Cl⁻ | bromide, Br⁻ | iodide, I⁻ |
|---|---|---|---|
| chlorine, Cl₂ | no reaction | reaction: Br₂ forms | reaction: I₂ forms |
| bromine, Br₂ | no reaction | no reaction | reaction: I₂ forms |
| iodine, I₂ | no reaction | no reaction | no reaction |
For chlorine and bromide ions: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq). Chlorine gains electrons and is reduced; bromide ions lose electrons and are oxidised.
Compare the elemental halogen with the halogen in the halide ion. Bromine displaces iodide but not chloride, because bromine is more reactive than iodine but less reactive than chlorine.
Predict properties of another Group VII element
Use the position of the unfamiliar halogen and the trends supplied by known Group VII elements to make bounded predictions.
| Position | Density | Reactivity | Likely appearance/state |
|---|---|---|---|
| above chlorine, such as fluorine | lower | higher | paler and gaseous |
| below iodine, such as astatine | higher | lower | darker and solid |
First place the element above, between or below known halogens. Extend each trend separately, and if numerical data are supplied, choose or estimate a value that continues the pattern without claiming unsupported precision.
Do not use only one trend to predict every property. Density increases and reactivity decreases down the group; state and colour predictions must be inferred from the specific information provided.
8.4 Transition elements
- Syllabus
- 0620–2026–2027
- Topic
- 8.4
- Level
- —
Recognise typical transition-element properties
Transition elements are metals in the central block of the Periodic Table. The required identifying pattern combines physical properties with characteristic compound and catalyst behaviour.
| Required property | What to recognise | Useful contrast with Group I metals |
|---|---|---|
| density | high | Group I metals have low densities |
| melting point | high | Group I metals have low melting points |
| compounds | often coloured | Group I compounds are usually white or colourless |
| catalytic behaviour | elements and their compounds often act as catalysts | Group I metals and compounds are not typically catalysts |
Iron, nickel, cobalt, copper, chromium and platinum are transition elements. Iron is used in the Haber process, nickel in hydrogenation, and transition-metal oxides can also catalyse reactions.
No single generic metal property is enough: conducting electricity or being malleable is common to many metals. Use the high-density/high-melting-point pattern together with coloured compounds or catalytic behaviour.
Use variable oxidation numbers
Transition elements can form ions with different oxidation numbers. The oxidation number is written as a Roman numeral in a compound's name.
| Ion | Oxidation number | Name in compounds | Example |
|---|---|---|---|
| Fe²⁺ | +2 | iron(II) | FeCl₂, iron(II) chloride |
| Fe³⁺ | +3 | iron(III) | FeCl₃, iron(III) chloride |
Find an unknown oxidation number by balancing charge. In Fe₂O₃, three O²⁻ ions total −6, so two iron ions total +6 and each iron is +3. In FeCl₂, two Cl⁻ ions require Fe²⁺.
The Roman numeral gives the oxidation number of the metal ion, not the number of metal atoms. Fe(III) means Fe³⁺; it does not mean that every formula contains three iron atoms.
8.5 Noble gases
- Syllabus
- 0620–2026–2027
- Topic
- 8.5
- Level
- —
Explain noble-gas unreactivity
The Group VIII noble gases are unreactive, monatomic gases. Monatomic means that their particles are single atoms, not bonded pairs or larger molecules.
| Noble gas | Electronic configuration | Outer shell |
|---|---|---|
| helium | 2 | first shell full with 2 electrons |
| neon | 2,8 | outer shell full with 8 electrons |
| argon | 2,8,8 | outer shell full with 8 electrons |
A full outer electron shell is already stable. Noble-gas atoms therefore have little tendency to gain, lose or share electrons, so they rarely form bonds and are very unreactive.
Do not say every noble gas has eight outer electrons: helium's only shell is full with two. Also, being a gas or monatomic does not cause inertness; the full outer shell explains it.