What you'll learn
This topic examines the vertical columns (groups) of the periodic table and how elements within each group share similar properties due to their electronic structure. Understanding group trends, reactivity patterns and characteristic reactions is essential for Paper 1 and appears frequently in 4-6 mark extended response questions. Questions commonly test your ability to predict properties, explain trends and write balanced equations for group reactions.
Key terms and definitions
Group — a vertical column in the periodic table containing elements with the same number of electrons in their outer shell
Alkali metals — the highly reactive metals in Group 1 that have one electron in their outer shell and form 1+ ions
Halogens — the reactive non-metals in Group 7 that have seven electrons in their outer shell and form 1- ions
Noble gases — the unreactive elements in Group 0 (or Group 8) with full outer electron shells
Displacement reaction — a reaction where a more reactive element takes the place of a less reactive element in a compound
Reactivity — the tendency of an element to undergo chemical reactions, determined by how easily atoms lose or gain electrons
Transition metals — the block of metallic elements between Groups 2 and 3 with characteristic properties including variable oxidation states and coloured compounds
Diatomic molecule — a molecule containing two atoms bonded together, such as Cl₂ or Br₂
Core concepts
Group 1: The alkali metals
Group 1 contains lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). These metals share characteristic properties:
- Soft metals that can be cut with a knife
- Low melting and boiling points compared to other metals
- Low density (lithium, sodium and potassium float on water)
- Form white ionic compounds that dissolve in water to give colourless solutions
- Always form 1+ ions in reactions because they lose their single outer electron
Reactivity trend: Reactivity increases down Group 1. Francium is the most reactive alkali metal.
Explanation: As you move down the group, the outer electron is further from the nucleus in an additional electron shell. The increased distance and increased shielding from inner electrons means the outer electron is less strongly attracted to the nucleus, making it easier to lose. This makes the atom more reactive.
Reactions with water:
All Group 1 metals react vigorously with water to produce a metal hydroxide and hydrogen gas:
- Lithium fizzes steadily, moves around the surface
- Sodium melts into a ball, fizzes vigorously, moves rapidly
- Potassium ignites with a lilac flame, may produce sparks, moves very rapidly
General equation: 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)
Example: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
The resulting solution is alkaline (pH 12-14) because of the metal hydroxide formed.
Reactions with oxygen: Group 1 metals tarnish rapidly in air as they react with oxygen. They are stored under oil to prevent this.
Example: 4Na(s) + O₂(g) → 2Na₂O(s)
Reactions with chlorine: Alkali metals react vigorously with chlorine gas to form white ionic chloride salts.
Example: 2Na(s) + Cl₂(g) → 2NaCl(s)
Group 7: The halogens
Group 7 contains fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂) and astatine (At₂). These elements exist as diatomic molecules and display a clear trend in physical properties:
- Fluorine: pale yellow gas (not studied at GCSE due to extreme reactivity)
- Chlorine: green gas at room temperature
- Bromine: red-brown liquid that vaporises easily to form brown gas
- Iodine: grey-black solid that sublimes to form purple vapour
Melting and boiling points increase down the group as the molecules become larger and intermolecular forces (Van der Waals forces) become stronger.
Reactivity trend: Reactivity decreases down Group 7. Fluorine is the most reactive halogen.
Explanation: Halogens react by gaining one electron to complete their outer shell and form a 1- ion. As you move down the group, the outer shell is further from the nucleus with more shielding from inner electrons. This means the attraction for an additional electron is weaker, making it harder to gain an electron. This makes the atom less reactive.
Displacement reactions: A more reactive halogen will displace a less reactive halogen from an aqueous solution of its salt. This provides evidence for the reactivity trend.
Chlorine displaces bromine and iodine:
Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
Observation: solution turns orange-brown as bromine forms
Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)
Observation: solution turns brown as iodine forms
Bromine displaces iodine but not chlorine:
Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)
Observation: solution turns brown as iodine forms
Iodine cannot displace chlorine or bromine because it is less reactive.
Reactions with alkali metals: Halogens react with Group 1 metals to form ionic halide salts. The reaction becomes less vigorous down the group as reactivity decreases.
Example: 2Na(s) + Cl₂(g) → 2NaCl(s)
Reactions with hydrogen: Halogens react with hydrogen gas to form hydrogen halides, which dissolve in water to form acidic solutions.
Example: H₂(g) + Cl₂(g) → 2HCl(g)
Hydrogen chloride dissolves in water to form hydrochloric acid:
HCl(g) + water → H⁺(aq) + Cl⁻(aq)
Group 0: The noble gases
Group 0 (sometimes called Group 8) contains helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). These elements are:
- Colourless gases at room temperature
- Monatomic (exist as single atoms, not molecules)
- Extremely unreactive (inert) with very few known compounds
- Non-flammable
Electronic structure explanation: Noble gases have full outer electron shells (helium has 2 electrons, all others have 8). This stable electronic arrangement means they do not need to lose, gain or share electrons, making them unreactive.
Physical property trends:
- Boiling points increase down the group as atoms become larger and intermolecular forces increase
- Density increases down the group
Uses (you may be tested on explaining uses based on properties):
- Helium: balloons and airships (low density, non-flammable)
- Neon: advertising signs (glows red-orange when electricity passes through it)
- Argon: filling light bulbs (inert atmosphere prevents the filament reacting), used in welding
- Krypton and Xenon: energy-efficient windows, lasers
Transition metals
The transition metals occupy the central block between Groups 2 and 3. Common examples include iron (Fe), copper (Cu), chromium (Cr), manganese (Mn), nickel (Ni) and zinc (Zn).
Characteristic properties:
- Good conductors of heat and electricity
- Hard, strong and dense metals
- High melting points (except mercury, which is liquid at room temperature)
- Much less reactive than Group 1 metals
- Form coloured compounds (e.g., copper compounds are often blue, iron(II) compounds are green, iron(III) compounds are orange-brown)
- Form ions with different charges (variable oxidation states), e.g., Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺
- Many are useful as catalysts (iron in the Haber process, nickel in hydrogenation of alkenes, manganese(IV) oxide in decomposition of hydrogen peroxide)
Comparison with Group 1 metals:
| Property | Group 1 | Transition metals |
|---|---|---|
| Reactivity | Very high | Low to moderate |
| Hardness | Soft | Hard |
| Density | Low | High |
| Melting point | Low | High |
| Ion charge | Always 1+ | Variable |
| Compounds | White | Often coloured |
Group patterns and the periodic table structure
Elements in the same group have:
- The same number of electrons in their outer shell
- Similar chemical properties
- Trends in physical properties down the group
- Trends in reactivity down the group
The group number tells you the number of outer shell electrons for Groups 1, 2 and 3, and for Groups 5, 6, 7 and 0 (where Group 0 elements have full outer shells).
Examples:
- Group 1 elements have 1 outer electron
- Group 7 elements have 7 outer electrons
- Group 0 elements have 8 outer electrons (except helium with 2)
Worked examples
Example 1: Predicting and explaining trends (6 marks)
Question: Rubidium is in Group 1, below potassium. Predict and explain what would be observed when rubidium is added to water. Include a balanced symbol equation in your answer.
Answer:
Rubidium would react very vigorously with water [1 mark], more vigorously than potassium [1 mark]. It would move very rapidly across the surface and would probably ignite immediately [1 mark], potentially exploding.
Rubidium is more reactive than potassium because the outer electron is further from the nucleus in an additional electron shell [1 mark]. The increased shielding means the outer electron is less strongly attracted to the nucleus and easier to lose [1 mark].
Balanced equation: 2Rb(s) + 2H₂O(l) → 2RbOH(aq) + H₂(g) [1 mark]
Example 2: Displacement reactions (4 marks)
Question: A student adds chlorine water to potassium iodide solution. Describe what the student would observe and explain why this reaction occurs. Include a balanced symbol equation.
Answer:
The colourless solution would turn brown [1 mark] as iodine is formed.
Chlorine is more reactive than iodine [1 mark], so it displaces iodine from the potassium iodide solution.
Balanced equation: Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq) [2 marks: 1 mark for correct formulae, 1 mark for balancing]
Example 3: Electronic structure and properties (3 marks)
Question: Explain why noble gases are unreactive.
Answer:
Noble gases have full outer electron shells [1 mark]. This electronic structure is very stable [1 mark], so they do not need to lose, gain or share electrons in chemical reactions [1 mark].
Common mistakes and how to avoid them
Mistake: Stating that reactivity increases down Group 7. Correction: Reactivity decreases down Group 7 as it becomes harder to gain an electron. Only metals (Groups 1 and 2) become more reactive down the group.
Mistake: Writing molecular formulae for halogens as single atoms (e.g., Cl instead of Cl₂). Correction: Halogens exist as diatomic molecules, so always write Cl₂, Br₂, I₂ in equations.
Mistake: Confusing which Group 1 metal is which during water reactions. Correction: Remember the reactivity order using LiSoKRubCaeFran (lithium, sodium, potassium, rubidium, caesium, francium). Potassium ignites with a lilac flame; sodium just melts and moves rapidly.
Mistake: Stating that distance from the nucleus is the only factor affecting reactivity. Correction: Always mention both increased distance AND increased shielding when explaining reactivity trends.
Mistake: Writing Group 0 as having 0 outer electrons. Correction: Group 0 elements have full outer shells (8 electrons for all except helium which has 2). The zero refers to their lack of reactivity.
Mistake: In displacement reactions, writing that a less reactive halogen displaces a more reactive one. Correction: Only a more reactive halogen can displace a less reactive halogen from its salt solution. Iodine cannot displace chlorine or bromine.
Exam technique for Groups in the Periodic Table
Command word 'Explain' requires you to give reasons. For reactivity trends, you must reference electron shells, distance from nucleus and shielding. Simply stating "more/less reactive" without explanation earns no marks.
Observation questions require you to state what you would see, not what is formed chemically. Write "the solution turns brown" rather than "iodine is produced" (though explaining what causes the observation may earn additional marks).
Balanced equations typically award 1 mark for correct formulae and 1 mark for balancing. State symbols may be required for full marks, especially in higher-tier questions.
Extended response questions (6 marks) on group trends require structured answers covering: the trend, observations/properties, the explanation using electronic structure, and often an equation. Plan your answer to ensure all elements are covered.
Quick revision summary
Group 1 metals become more reactive down the group as the outer electron is easier to lose. They react with water to form alkaline solutions and hydrogen. Group 7 halogens become less reactive down the group as gaining an electron becomes harder. More reactive halogens displace less reactive ones. Noble gases (Group 0) are unreactive due to full outer shells. Transition metals have high melting points, form coloured compounds, show variable oxidation states and act as catalysts. Group number indicates outer shell electrons for most groups.