Part 1: Group 2 reactions and solubility

Use the ns² outer configuration to explain electron loss, then distinguish metal redox reactions from acid–base reactions of oxides and hydroxides.

Why magnesium to barium becomes more reactive

Group 2 atoms have two outer electrons and commonly form M²⁺ ions by losing both. Down the group, the outer electrons occupy shells farther from the nucleus and are more shielded by inner electrons. These effects outweigh increasing proton number, so the first and second ionisation energies generally decrease and electron loss becomes easier.

This explains the broad increase in reactivity from magnesium through calcium and strontium to barium. A measured reaction rate also depends on temperature, surface area and surface films; it is not determined by ionisation energy alone. Magnesium’s oxide coating can slow its initial reaction, so clean comparable pieces before a controlled comparison.

When discussing ionisation energy, refer to attraction between the nucleus and the outer electron. When discussing the metal’s melting temperature, refer to metallic bonding. Group 2 melting data contain structural irregularities and are not a simple monotonic reactivity scale.

Reactions with oxygen and chlorine

Group 2 metals are oxidised from 0 to +2 when forming ordinary oxides or chlorides. Oxygen is reduced from 0 to −2 in an oxide; chlorine is reduced from 0 to −1 in a chloride. Magnesium burns with a brilliant white flame to give white MgO. The general simple-oxide equation is 2M + O₂ → 2MO; a chloride forms by M + Cl₂ → MCl₂.

For calcium, 2Ca + O₂ → 2CaO and Ca + Cl₂ → CaCl₂. The analogous Mg, Sr and Ba chloride formulae are MgCl₂, SrCl₂ and BaCl₂. Count the diatomic chlorine molecule before balancing. Conditions can affect oxygen products: barium readily forms a peroxide in excess oxygen, Ba + O₂ → BaO₂, where oxygen is −1. Use the oxide or peroxide formula specified by the question rather than silently changing it.

These are vigorous oxidation reactions. In practical work, use the prescribed small-scale demonstration and eye protection; do not stare at burning magnesium, and handle chlorine only with appropriate ventilation and supervision. The white light from burning magnesium metal is not a magnesium-ion flame-test colour.

Steam and liquid water do not give identical magnesium products

Magnesium reacts very slowly with cold water; warmer water can produce magnesium hydroxide and hydrogen. With steam at high temperature, the main AS product is magnesium oxide and hydrogen. State the conditions because the equations differ.

Calcium, strontium and barium react with cold water with increasing vigour to produce the metal hydroxide and hydrogen. Fizzing shows gas evolution, the metal is consumed and the mixture warms. Calcium may give a cloudy suspension because calcium hydroxide is only moderately soluble; the hydroxides become more soluble down the group.

In a comparison, keep temperature, exposed metal surface area and water volume controlled and record a measurable feature such as gas volume in a fixed time. Small pieces and a shield may be appropriate for more reactive metals under the teacher’s risk assessment. Hydrogen is flammable, so keep the reaction away from ignition sources.

Basic oxides and hydroxides neutralise acids

A Group 2 oxide reacts with water to form its hydroxide: MO + H₂O → M(OH)₂. Calcium oxide reacts exothermically in the slaking process. Magnesium oxide reacts less readily and gives sparingly soluble magnesium hydroxide. An alkaline solution contains dissolved OH⁻ ions; a suspension may contain much more undissolved hydroxide than dissolved material.

Oxides and hydroxides react with dilute acid to form a salt and water. These are acid–base reactions with no change in the metal’s +2 oxidation number. They do not normally produce hydrogen gas; the metal itself can produce hydrogen when oxidised by acid.

For an insoluble hydroxide, a useful ionic equation is Mg(OH)₂(s) + 2H⁺(aq) → Mg²⁺(aq) + 2H₂O(l). For dissolved hydroxide, the core neutralisation equation is H⁺ + OH⁻ → H₂O. Use states consistent with whether the reagent is a solid suspension or a solution.

Hydroxides become more soluble; sulfates become less soluble

Do not reverse these two trends. Aqueous Mg²⁺ with OH⁻ gives a white Mg(OH)₂ precipitate. Aqueous Ba²⁺ with sulfate gives white BaSO₄. ‘Insoluble’ in qualitative analysis usually means sufficiently low solubility to form a visible solid, not absolutely zero ions dissolved.

A sulfate coating can slow reaction of a metal or oxide with sulfuric acid. For example, insoluble BaSO₄ may block contact between barium and acid, even though barium is intrinsically very reactive. This is a physical barrier, not evidence that barium has become less reactive than magnesium.

Worked supplied-data calculation: a hydroxide M(OH)₂ with molar mass 74.1 g mol⁻¹ has a stated solubility of 0.148 g per 100 cm³ of solution at a fixed temperature. Dissolved amount = 0.148 ÷ 74.1 = 0.001997 mol. Volume = 0.100 dm³, so concentration of M²⁺ = 0.01997 mol dm⁻³. Each formula unit gives two OH⁻, so [OH⁻] = 0.0399 mol dm⁻³ to three significant figures. Equal masses of different hydroxides do not automatically give equal hydroxide-ion concentrations.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.

Sources