Chemistry of the Elements

What you'll learn

This revision guide covers the essential chemistry of elements in Groups 1, 7 and 0 of the periodic table, transition metals, and the reactivity series. You'll understand trends in reactivity, physical properties, and reactions of key elements tested in the Pearson Edexcel International IGCSE Chemistry examination. These topics form a significant portion of Paper 1 and Paper 2 assessments.

Key terms and definitions

Alkali metals — elements in Group 1 of the periodic table; soft, reactive metals that form alkaline solutions when reacting with water

Halogens — elements in Group 7 of the periodic table; reactive non-metals that form diatomic molecules and exist in different states at room temperature

Displacement reaction — a reaction where a more reactive element takes the place of a less reactive element in a compound

Noble gases — elements in Group 0 (or Group 18) of the periodic table; unreactive gases with full outer electron shells

Transition metals — elements found in the central block of the periodic table between Groups 2 and 3; they have characteristic properties including variable oxidation states and coloured compounds

Reactivity series — an arrangement of metals in order of their reactivity, from most reactive (potassium) to least reactive (platinum)

Diatomic molecule — a molecule containing two atoms chemically bonded together, such as Cl₂, Br₂ or I₂

Halide — a compound formed when a halogen reacts with another element, typically containing a halide ion (Cl⁻, Br⁻ or I⁻)

Core concepts

Group 1: The alkali metals

The alkali metals (lithium, sodium, potassium, rubidium, caesium and francium) occupy Group 1 of the periodic table. Only lithium, sodium and potassium are commonly examined at IGCSE level.

Physical properties and trends:

• 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
• Good conductors of heat and electricity

Going down Group 1:

• Melting points decrease
• Boiling points decrease
• Density increases
• Reactivity increases

Chemical reactions:

Reaction with water produces a metal hydroxide and hydrogen gas:

• 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
• 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Observations:

• Lithium: fizzes steadily, moves slowly across the water surface
• Sodium: melts into a ball, moves quickly across the water, may ignite with an orange flame
• Potassium: melts immediately, moves very rapidly, burns with a lilac flame

Reaction with oxygen produces metal oxides. Lithium forms lithium oxide:

• 4Li(s) + O₂(g) → 2Li₂O(s)

Sodium and potassium may also form peroxides.

Reaction with chlorine produces white ionic salts (metal chlorides):

• 2Na(s) + Cl₂(g) → 2NaCl(s)

All Group 1 metals form ionic compounds with 1+ ions (Li⁺, Na⁺, K⁺). These compounds dissolve in water to form alkaline solutions (pH 12-14).

Explaining the reactivity trend:

Group 1 metals react by losing their single outer electron to form 1+ ions. Going down the group:

• Atomic radius increases
• Outer electron is further from the nucleus
• Attraction between nucleus and outer electron decreases
• Outer electron is lost more easily
• Therefore reactivity increases

Group 7: The halogens

The halogens (fluorine, chlorine, bromine, iodine and astatine) occupy Group 7 of the periodic table. Fluorine and astatine are rarely examined at IGCSE level.

Physical properties and trends:

At room temperature (25°C):

• Chlorine (Cl₂): pale green gas
• Bromine (Br₂): red-brown liquid, produces brown vapour
• Iodine (I₂): grey-black solid, produces purple vapour when heated (sublimation)

All exist as diatomic molecules (two atoms bonded by a covalent bond).

Going down Group 7:

• Melting points increase
• Boiling points increase
• Colour becomes darker
• Reactivity decreases

Chemical reactions:

Halogens react with metals to form ionic halides:

• 2Fe(s) + 3Cl₂(g) → 2FeCl₃(s)
• 2Na(s) + Br₂(l) → 2NaBr(s)

Halogens react with hydrogen to form hydrogen halides:

• H₂(g) + Cl₂(g) → 2HCl(g)

These hydrogen halides dissolve in water to form acidic solutions:

• HCl(g) + H₂O(l) → HCl(aq) [hydrochloric acid]

Displacement reactions:

A more reactive halogen will displace a less reactive halogen from an aqueous solution of its halide. This demonstrates the reactivity order: Cl₂ > Br₂ > I₂.

Examples:

• Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
• Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)
• Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)

But:

• Br₂(aq) + KCl(aq) → no reaction (bromine is less reactive than chlorine)
• I₂(aq) + KBr(aq) → no reaction (iodine is less reactive than bromine)

Explaining the reactivity trend:

Halogens react by gaining one electron to complete their outer shell and form 1- ions. Going down the group:

• Atomic radius increases
• Outer shell is further from the nucleus
• Attraction between nucleus and incoming electron decreases
• Electron is gained less easily
• Therefore reactivity decreases

Group 0: The noble gases

The noble gases (helium, neon, argon, krypton, xenon and radon) occupy Group 0 (or Group 18) of the periodic table.

Properties:

• Colourless gases at room temperature
• Exist as single atoms (monatomic)
• Very unreactive (inert) because they have full outer electron shells
• Non-flammable

Going down Group 0:

• Boiling points increase
• Density increases

Uses:

• Helium: party balloons, airships (low density and non-flammable)
• Neon: advertising signs and lights (glows red-orange when electricity passes through)
• Argon: light bulbs, welding (provides inert atmosphere to prevent reaction)

Transition metals

The transition metals are located in the central block of the periodic table between Groups 2 and 3. Common examples include iron, copper, chromium, manganese, nickel and zinc.

Properties:

• High melting points (except mercury, which is liquid)
• High density
• Strong and hard
• Good conductors of heat and electricity
• Less reactive than Group 1 metals
• Form coloured compounds
• Have catalytic properties
• Show variable oxidation states (e.g., iron forms Fe²⁺ and Fe³⁺)

Examples of coloured compounds:

• Copper(II) compounds: blue (e.g., CuSO₄ solution)
• Iron(II) compounds: pale green (e.g., FeSO₄ solution)
• Iron(III) compounds: yellow/brown (e.g., FeCl₃ solution)
• Chromium(III) compounds: green (e.g., Cr₂O₃)

Catalytic examples:

• Iron in the Haber process (N₂ + 3H₂ ⇌ 2NH₃)
• Vanadium(V) oxide in the Contact process (2SO₂ + O₂ ⇌ 2SO₃)
• Nickel in hydrogenation of alkenes
• Manganese(IV) oxide in decomposition of hydrogen peroxide

Reactivity series of metals

The reactivity series arranges metals in order of decreasing reactivity:

Potassium (most reactive) Sodium Calcium Magnesium Aluminium Carbon (non-metal, included for extraction purposes) Zinc Iron Tin Lead Hydrogen (non-metal, included for reference) Copper Silver Gold Platinum (least reactive)

Applications of the reactivity series:

1. Predicting reactions with water

   • Metals above hydrogen react with water or steam
   • More reactive metals react more vigorously
   • Potassium, sodium and calcium react with cold water
   • Magnesium reacts slowly with cold water, rapidly with steam
   • Zinc and iron react with steam only

2. Predicting reactions with acids

   • Metals above hydrogen react with dilute acids to produce a salt and hydrogen gas
   • Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)
   • Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)
   • Metals below hydrogen do not react with dilute acids

3. Displacement reactions

   • A more reactive metal displaces a less reactive metal from a solution of its salt
   • Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)
   • Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
   • Cu(s) + MgSO₄(aq) → no reaction (copper is less reactive than magnesium)

4. Extraction of metals from ores

   • Metals above carbon must be extracted by electrolysis (expensive)
   • Metals below carbon can be extracted by reduction with carbon or carbon monoxide (cheaper)
   • Example: iron extraction in a blast furnace using carbon monoxide

Tests for metal ions and halide ions

Flame tests for metal ions:

• Lithium (Li⁺): crimson/red flame
• Sodium (Na⁺): yellow/orange flame
• Potassium (K⁺): lilac flame
• Calcium (Ca²⁺): orange-red flame
• Copper (Cu²⁺): blue-green flame

Test for halide ions:

Add dilute nitric acid followed by silver nitrate solution:

• Chloride (Cl⁻): white precipitate of silver chloride
• Bromide (Br⁻): cream precipitate of silver bromide
• Iodide (I⁻): yellow precipitate of silver iodide

The nitric acid removes any carbonate ions that would interfere by producing a white precipitate.

Worked examples

Example 1: Describe and explain what you would observe when a small piece of sodium is added to a large beaker of water. Write a balanced equation for the reaction. (4 marks)

Solution:

• Sodium floats on the water / moves across the surface (1 mark)
• Fizzing / effervescence observed (due to hydrogen gas production) (1 mark)
• Sodium may melt into a ball / may catch fire with an orange flame (1 mark)
• 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) (1 mark)

Example 2: A student adds chlorine water to separate solutions of potassium bromide and potassium iodide. (a) State what the student would observe in each case. (2 marks) (b) Explain why chlorine can displace both bromine and iodine from their compounds. (2 marks) (c) Write a balanced ionic equation for the reaction between chlorine and potassium iodide solution. (2 marks)

Solution: (a)

• With potassium bromide: solution turns orange/brown (due to bromine) (1 mark)
• With potassium iodide: solution turns brown/black (due to iodine) (1 mark)

(b)

• Chlorine is more reactive than both bromine and iodine (1 mark)
• Therefore chlorine displaces both bromine and iodine from their compounds (1 mark)

Alternative: Chlorine gains electrons more easily than bromine or iodine (because outer shell is closer to nucleus / stronger nuclear attraction) (2 marks)

(c) Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq) (2 marks — award 1 mark for correct species, 1 mark for balancing)

Example 3: Use the reactivity series to predict whether the following reactions will occur. If a reaction occurs, write a balanced equation. (4 marks) (a) Zinc metal with copper(II) sulfate solution (b) Copper metal with magnesium sulfate solution

Solution: (a) Reaction occurs because zinc is more reactive than copper (1 mark) Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) (1 mark)

(b) No reaction occurs because copper is less reactive than magnesium / magnesium is more reactive than copper (1 mark) Cu(s) + MgSO₄(aq) → no reaction (1 mark)

Common mistakes and how to avoid them

• Confusing reactivity trends: Remember Group 1 reactivity increases down the group but Group 7 reactivity decreases down the group. This is because Group 1 metals lose electrons (easier when further from nucleus) while Group 7 non-metals gain electrons (harder when further from nucleus).

• Incorrect halogen displacement predictions: Always check the reactivity order (Cl > Br > I). A halogen can only displace halogens below it in the group, not above.

• Forgetting diatomic molecules: Halogens exist as Cl₂, Br₂ and I₂, not as single atoms. Always write them with a subscript 2 in equations.

• Misinterpreting observations: "Fizzing" and "effervescence" both indicate gas production. Be specific about which gas (hydrogen when metals react with water or acids). Use precise colour descriptions: "orange-red" not just "red", "lilac" not "purple" for potassium flame.

• Incomplete flame test method: Always state that the metal compound should be placed on a nichrome wire that has been cleaned with dilute hydrochloric acid and heated until no colour appears. This removes contamination.

• Forgetting to add nitric acid in halide tests: The nitric acid is essential to prevent false positives from carbonate ions. Always mention adding dilute nitric acid before adding silver nitrate solution.

Exam technique for "Chemistry of the Elements"

• Command words matter: "State" requires a brief answer without explanation (1 mark). "Explain" requires a reason and often earns 2 marks. "Describe" requires observation details. "Predict" requires you to apply patterns or rules to new situations.

• Reactivity series questions: When comparing metals or halogens, always refer explicitly to the reactivity series or state which element is "more reactive" or "less reactive". Vague answers like "because it's higher up" lose marks without naming the elements.

• Balanced equations: Check coefficients carefully. Equation questions typically award 1 mark for correct formulae and 1 mark for correct balancing. State symbols may be required for full marks if the question asks for them.

• Extended response questions: Structure answers logically. For 4-6 mark questions on reactivity trends, describe the observation, state the pattern/trend, then explain using electron structure. Use scientific vocabulary: "atomic radius", "nuclear attraction", "outer shell".

Quick revision summary

Group 1 metals are soft, reactive and form alkaline solutions; reactivity increases down the group as the outer electron is lost more easily. Group 7 halogens exist as diatomic molecules; reactivity decreases down the group as electrons are gained less easily. More reactive halogens displace less reactive ones. Noble gases are unreactive due to full outer shells. Transition metals form coloured compounds, act as catalysts and show variable oxidation states. The reactivity series predicts displacement reactions, reactions with acids/water, and extraction methods. Metal ions produce characteristic flame colours, and halides form coloured precipitates with silver nitrate.