What you'll learn
The Periodic Table forms the foundation for understanding element behaviour and chemical reactions in CIE IGCSE Chemistry. Questions on this topic appear consistently across Paper 2, Paper 4 and Paper 6, testing your knowledge of element arrangement, group properties, periodic trends and the ability to predict chemical behaviour. Mastery of this topic is essential for success in both core and extended tier examinations.
Key terms and definitions
Periodic Table — a tabular arrangement of chemical elements organised by increasing atomic (proton) number, showing recurring patterns in chemical and physical properties.
Group — a vertical column in the Periodic Table containing elements with the same number of outer shell (valence) electrons, resulting in similar chemical properties.
Period — a horizontal row in the Periodic Table, representing elements whose atoms have the same number of electron shells.
Atomic number — the number of protons in the nucleus of an atom, which defines the element and determines its position in the Periodic Table.
Metallic character — the tendency of an element to lose electrons and form positive ions, characteristic of metals.
Transition elements — metallic elements found in the central block between Groups II and III, characterised by variable oxidation states and coloured compounds.
Noble gases — Group VIII (or Group 0) elements with complete outer electron shells, making them extremely unreactive.
Alkali metals — Group I elements that are highly reactive metals, forming alkaline solutions when they react with water.
Core concepts
Structure and organisation of the Periodic Table
The modern Periodic Table arranges all known elements by atomic number (number of protons) from left to right and top to bottom. This organisation reveals periodic patterns in properties.
Key structural features:
- Groups numbered I to VIII (or 1-18 in IUPAC notation) run vertically
- Periods numbered 1 to 7 run horizontally
- Elements in the same group have identical numbers of outer shell electrons
- Period number indicates the number of electron shells in an atom
- The zigzag line separates metals (left) from non-metals (right)
- Metalloids like silicon and germanium lie along this boundary
For CIE IGCSE examinations, focus on the first 36 elements (hydrogen to krypton) and be able to identify their positions and properties.
Group I: The alkali metals
Group I elements (lithium, sodium, potassium, rubidium, caesium, francium) share characteristic properties due to having one electron in their outer shell.
Physical 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)
- Shiny when freshly cut, but tarnish rapidly in air
Chemical properties:
- Extremely reactive metals that must be stored under oil
- React vigorously with water to produce hydrogen gas and metal hydroxides
- Form ionic compounds with non-metals
- Always form +1 ions (lose their single outer electron)
- React with oxygen to form metal oxides
Example reaction with water:
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
Trend down Group I:
- Reactivity increases (lithium < sodium < potassium)
- Melting points decrease
- Atoms get larger (atomic radius increases)
- Outer electron is further from the nucleus, so more easily lost
Group VII: The halogens
Group VII elements (fluorine, chlorine, bromine, iodine, astatine) have seven electrons in their outer shell and exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).
Physical properties at room temperature:
- Fluorine: pale yellow gas
- Chlorine: green-yellow gas
- Bromine: red-brown liquid (only liquid non-metal)
- Iodine: grey-black solid that sublimates to purple vapour
Chemical properties:
- Reactive non-metals that form -1 ions (gain one electron)
- React with metals to form ionic salts (metal halides)
- React with hydrogen to form hydrogen halides (HCl, HBr, HI)
- Form covalent compounds with other non-metals
Trend down Group VII:
- Reactivity decreases (fluorine > chlorine > bromine > iodine)
- Melting and boiling points increase
- Colour intensity increases
- Atoms get larger (atomic radius increases)
- Outer shell is further from nucleus, so harder to attract an additional electron
Displacement reactions:
A more reactive halogen displaces a less reactive halogen from its compound:
Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
Chlorine displaces bromine because chlorine is more reactive. This reaction produces an orange solution as bromine is formed.
Group VIII (Group 0): The noble gases
Group VIII elements (helium, neon, argon, krypton, xenon, radon) are characterised by complete outer electron shells, making them uniquely stable.
Properties:
- Colourless, odourless gases at room temperature
- Extremely unreactive (inert) due to stable electron configuration
- Do not form molecules (exist as single atoms: monatomic)
- Do not readily form compounds
- Low boiling points
Trend down Group VIII:
- Boiling points increase as atomic mass increases
- Density increases
Uses based on inertness:
- Helium: filling balloons, airships (low density, non-flammable)
- Neon: advertising signs (glows red-orange when electricity passes through)
- Argon: inert atmosphere in light bulbs and welding
Transition elements
The transition elements occupy the central block between Groups II and III (scandium to zinc in Period 4).
Characteristic properties:
- Hard, strong metals with high melting points
- Good conductors of heat and electricity
- High density compared to Group I metals
- Less reactive than Group I metals
- Form coloured compounds (e.g. copper sulfate is blue, iron(II) sulfate is green)
- Often act as catalysts (iron in Haber process, nickel in hydrogenation)
- Show variable oxidation states (e.g. iron forms Fe²⁺ and Fe³⁺)
- Form complex ions with ligands
Common examples for CIE IGCSE:
- Iron (Fe): used in construction, forms rust
- Copper (Cu): electrical wiring, water pipes
- Zinc (Zn): galvanising steel
- Silver (Ag): jewellery, mirrors
- Gold (Au): jewellery, electronics
Periodic trends across periods
Across Period 3 (sodium to argon):
Metallic to non-metallic character:
- Sodium, magnesium and aluminium are metals
- Silicon is a metalloid
- Phosphorus, sulfur, chlorine and argon are non-metals
Atomic structure changes:
- All Period 3 elements have three electron shells
- Atomic number increases by one across the period
- Number of outer shell electrons increases from 1 to 8
Properties trend:
- Metallic character decreases across the period
- Elements change from metals (electron donors) to non-metals (electron acceptors)
- Melting points generally increase from Na to Si, then decrease
- Silicon has highest melting point due to giant covalent structure
Oxide acidity trend:
- Metal oxides (Na₂O, MgO, Al₂O₃) are basic or amphoteric
- Non-metal oxides (P₄O₁₀, SO₂, SO₃) are acidic
- This represents a transition from basic to acidic character
Example reactions:
Na₂O(s) + H₂O(l) → 2NaOH(aq) — basic oxide forms alkaline solution
SO₂(g) + H₂O(l) → H₂SO₃(aq) — acidic oxide forms acidic solution
Predicting properties from position
The Periodic Table allows prediction of element behaviour:
From group number:
- Number of outer electrons equals group number (for Groups I-VII)
- Charge of common ion: Group I = +1, Group II = +2, Group VII = -1
- Chemical reactivity patterns within groups
From period number:
- Number of electron shells equals period number
- Helps predict atomic size
Metal or non-metal:
- Position relative to the zigzag line determines bonding type
- Metals form ionic compounds with non-metals
- Non-metals form covalent compounds with other non-metals
Worked examples
Example 1: Group trends
Question: Explain why potassium is more reactive than sodium. Your answer should refer to electronic structure. [3 marks]
Answer:
- Both potassium and sodium have one electron in their outer shell [1]
- Potassium has four electron shells while sodium has three electron shells / potassium atom is larger [1]
- The outer electron in potassium is further from the nucleus, so it is lost more easily / experiences less attraction from the nucleus [1]
Example 2: Displacement reactions
Question: A student adds chlorine water to a solution of potassium iodide.
(a) State what the student would observe. [1]
(b) Write a balanced symbol equation for the reaction. [2]
(c) Explain why this reaction occurs. [2]
Answer:
(a) The colourless solution turns brown/yellow [1]
(b) Cl₂ + 2KI → 2KCl + I₂ [2] (1 mark for correct formulae, 1 mark for balancing)
(c) Chlorine is more reactive than iodine [1], so chlorine displaces iodine from its compound / chlorine is a stronger oxidising agent [1]
Example 3: Periodic Table interpretation
Question: An element X is in Group VII and Period 3 of the Periodic Table.
(a) How many electron shells does an atom of element X have? [1]
(b) How many electrons are in the outer shell of an atom of element X? [1]
(c) Identify element X. [1]
(d) Predict the formula of the compound formed between X and magnesium. [1]
Answer:
(a) 3 [1] (Period number = number of shells)
(b) 7 [1] (Group number = outer electrons for Groups I-VII)
(c) Chlorine / Cl [1]
(d) MgCl₂ [1] (Mg²⁺ and Cl⁻ ions)
Common mistakes and how to avoid them
Mistake: Confusing groups with periods. Students write "Group 3" when they mean "Period 3". Correction: Groups are vertical columns (same outer electrons), periods are horizontal rows (same number of shells). Remember: groups go UP and DOWN like a GROUP of people standing.
Mistake: Stating that reactivity increases down Group VII, when it actually decreases. Correction: Halogens become less reactive down the group because they need to gain an electron, and larger atoms attract additional electrons less effectively. This is opposite to Group I metals which become more reactive.
Mistake: Writing that noble gases have eight outer electrons (helium only has two). Correction: Noble gases have full outer shells: helium has 2 electrons (first shell full), all others have 8 (second and subsequent shells full).
Mistake: Claiming transition metals are between Group II and Group III rather than identifying their actual position. Correction: Transition elements occupy the block from Group III to Group II in modern tables (scandium to zinc). For CIE IGCSE, focus on common examples: iron, copper, zinc, silver, gold.
Mistake: Forgetting that halogens exist as diatomic molecules (Cl₂, Br₂, I₂) in equations. Correction: Always write halogen elements as X₂ when writing equations. Single atoms only exist in ionic compounds as halide ions (Cl⁻, Br⁻, I⁻).
Mistake: Stating that atomic number is the number of electrons rather than protons. Correction: Atomic number always equals the number of protons. In a neutral atom, this also equals the number of electrons, but ions have different numbers of electrons.
Exam technique for The Periodic Table
Explain questions about reactivity trends require reference to atomic structure (number of shells, distance of outer electrons from nucleus, ease of losing/gaining electrons). Simply stating which is more reactive scores no marks without explanation.
Prediction questions test your ability to apply group trends. If asked about an unfamiliar element, identify its group position and apply known patterns. For example, if asked about rubidium, recognise it is below potassium in Group I and predict higher reactivity and lower melting point.
Balanced equations for halogen displacement reactions are common. Remember: more reactive halogen + less reactive halide compound → more reactive halide compound + less reactive halogen. Check all halogens are written as X₂ when elements.
Data-based questions may provide information about an unknown element and ask you to predict its group. Use clues: charge of ion indicates group, physical state suggests position in Group VII, reactivity data indicates relative position in a group.
Quick revision summary
The Periodic Table arranges elements by atomic number into groups (vertical, same outer electrons) and periods (horizontal, same number of shells). Group I metals increase in reactivity down the group as outer electrons are more easily lost. Group VII halogens decrease in reactivity down the group as atoms attract additional electrons less effectively. Group VIII noble gases are unreactive due to full outer shells. Transition elements show variable oxidation states and form coloured compounds. Across periods, elements change from metals to non-metals, and oxides change from basic to acidic. Position predicts properties and reactivity.