The periodic table

Use H2 Chemistry periodic trends to explain Period 3 properties and halogen reactions, then practise deductions from unfamiliar chemical data.

GCE A-Level H2 Chemistry 2027
6 lessons

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Before you begin

This topic explains the patterns in the periodic table. It starts with atomic radius, ionisation energy and electronegativity, uses them to explain the structures and compounds of the Period 3 elements and the trends down Groups 2 and 17, and ends by deducing an unknown element from data.

Be comfortable with: electron configuration, first ionisation energy trends and chemical bonding. The Group 2 and Group 17 lessons introduce the electrode potentials they use.

Learning goals

Periodic Trends (Radius, IE, EN)
Period 3 Elements: Melting Point and Conductivity
Period 3 Oxides and Chlorides
Group 2 Chemistry Trends
Group 17 Chemistry Trends
Deducing Unknown Elements from Data

Syllabus statements covered

recognise variation in the electronic configurations across a Period and down a Group
describe and explain qualitatively the general trends and variations in atomic radius, ionic radius, first ionisation energy and electronegativity: — across a Period in terms of shielding and nuclear charge
describe and explain qualitatively the general trends and variations in atomic radius, ionic radius, first ionisation energy and electronegativity: — down a Group in terms of increasing number of electron shells, shielding and nuclear charge
interpret the variation in melting point and in electrical conductivity across a Period in terms of structure and bonding in the elements (metallic, giant molecular, or simple molecular)
state and explain the variation in the highest oxidation number of the elements in oxides (for Na2O; MgO; Al2O3; SiO2; P4O10; SO3) and chlorides (for NaCl; MgCl2; AlCl3; SiCl4; PCl5)
state and explain the variation in bonding in oxides and chlorides in terms of electronegativity (with the exception of AlCl3)
describe the reactions of the oxides with water (for Na2O; MgO; Al2O3; SiO2; P4O10; SO3)
describe and explain the acid/base behaviour of oxides (for Na2O; MgO; Al2O3; SiO2; P4O10; SO3) and hydroxides (for NaOH; Mg(OH)2; Al(OH)3), including, where relevant, amphoteric behaviour in reaction with sodium hydroxide (only) and acids
describe and explain the reactions of the chlorides with water (for NaCl; MgCl2; AlCl3; SiCl4; PCl5)
suggest the types of structure and bonding present in the oxides and chlorides from observations of their chemical and physical properties
describe and deduce from E⦵ values the relative reactivity of elements of: — Group 2 as reducing agents
describe and explain the trend in thermal stability of: — Group 2 carbonates in terms of the charge density of the cation and the polarisability of the large anion
describe and explain the trend in volatility of the Group 17 elements in terms of instantaneous dipole-induced dipole attraction
describe and deduce from E⦵ values the relative reactivity of elements of: — Group 17 as oxidising agents
describe and explain the trend in thermal stability of: — Group 17 hydrides in terms of bond energies
predict the characteristic properties of an element in a given Group by using knowledge of chemical periodicity
deduce the nature, possible position in the Periodic Table, and identity of unknown elements from given information of physical and chemical properties

Lessons

Work through them in order.

Practise and check

Or choose

Topic reference

Quick reference

Trend	Direction	Cause
atomic radius across a period	decreases	nuclear charge rises; shielding changes little
atomic radius down a group	increases	more shells; outer electrons further away and more shielded
Period 3 melting point	rises Na to Si, then falls sharply at P	metallic, then giant covalent, then simple molecular
Period 3 oxides	basic, then amphoteric (Al₂O₃), then acidic	bonding changes from ionic to covalent
Group 2 metals as reducing agents	stronger down the group	E⦵(M²⁺/M) more negative
Group 2 carbonate thermal stability	increases down the group	larger cation polarises the carbonate ion less
Group 17 volatility	decreases down the group	more electrons; stronger London forces
Group 17 oxidising power	decreases down the group	E⦵(X₂/X⁻) less positive
hydrogen halide thermal stability	decreases from HCl to HI	H–X bond longer and weaker

Core knowledge to remember

Across a period, nuclear charge increases while shielding changes little, so the outer electrons are held more strongly.
Down a group, the outer electrons are further from the nucleus and more shielded, so they are held less strongly.
In an isoelectronic series, the ion with the higher nuclear charge is smaller.
The highest oxidation number across Period 3 rises from +1 in Na₂O and NaCl to +6 in SO₃; the chlorides stop at +5 in PCl₅.
Al₂O₃ and Al(OH)₃ are amphoteric. AlCl₃ is mainly covalent because Al³⁺ strongly polarises the chloride ion.
Covalent chlorides are hydrolysed by water to give HCl; ionic chlorides only dissolve.
E⦵ values in data tables are for reductions as written.
A halogen displaces the halide ions of the halogens below it.

Common mistakes

Stating a trend without its cause: nuclear charge, shielding and distance, and the attraction that results.
Treating shielding as the main cause of the trend across a period.
Explaining the melting point of a simple molecular element with covalent bonds instead of forces between molecules.
Calling an oxide amphoteric without showing a reaction with an acid and with an alkali.
Reading a more negative E⦵(M²⁺/M) as a weaker reducing metal, or a less positive E⦵(X₂/X⁻) as a stronger oxidising halogen.
Explaining hydrogen-halide stability with forces between molecules instead of the H–X bond.
Writing halogens as single atoms, such as Cl for Cl₂.