Group 2 — OCR A-Level Chemistry

Group 2 metals (Mg, Ca, Sr, Ba) react by losing two outer electrons to form 2⁺ ions, so they act as reducing agents and are oxidised from oxidation state 0 to 2⁺. Reactivity increases down the group because the first and second ionisation energies decrease as atomic radius increases and shielding grows, so the outer electrons are lost more readily. They react with oxygen to form oxides, with water to form hydroxides and hydrogen, and with dilute acids to form salts and hydrogen. In each case the metal is oxidised and another species is reduced, for example hydrogen in water or acid is reduced from 1⁺ to 0. Solubility of hydroxides increases down the group while solubility of sulfates decreases.

Group 2 elements (Mg, Ca, Sr, Ba) each have two electrons in their outer s subshell, written as an s² configuration such as [Ne] 3s² for magnesium or [Ar] 4s² for calcium. These two outer electrons are relatively easily removed because the resulting 2⁺ ion has a stable noble-gas configuration and the ionisation energies, though endothermic, are achievable in reactions. When a Group 2 atom reacts, it loses both outer s electrons and is itself oxidised, while another species gains electrons and is reduced. The product is a 2⁺ cation, for example Mg²⁺ or Ca²⁺, so the group consistently forms compounds in which the metal has oxidation state +2. Recognising this pattern lets you write formulae such as MgCl₂ and CaO and balance redox half-equations showing the metal losing two electrons.

Reactivity of the Group 2 metals increases down the group from magnesium to barium. Each metal is oxidised, losing its two outer s electrons to form a 2⁺ ion, while another species is reduced. As you descend, the atoms get larger and the outer electrons are further from the nucleus and more shielded, so the first and second ionisation energies fall and the metal loses its electrons more readily. This trend is shown by the metals' redox reactions, for example with water, with oxygen and with dilute acid, where barium reacts most vigorously and magnesium least. Comparing the vigour and the ease of forming the 2⁺ ion lets you rank the metals and explain the trend in terms of ionisation energy and atomic radius.

This item sits in the Group 2 section, where you study how the alkaline earth metals magnesium, calcium, strontium and barium react and how their compounds behave. The prompt names oxygen as the reagent, so the task is to identify what a Group 2 metal forms when it is heated in oxygen. Group 2 metals are oxidised here: each metal atom loses two electrons to form a 2⁺ ion, and oxygen gains electrons to form the oxide ion O²⁻. The product is a solid metal oxide with the general formula MO, for example magnesium oxide MgO or calcium oxide CaO. Write the balanced equation with state symbols, such as 2Mg(s) + O₂(g) → 2MgO(s), and note that the metal is oxidised while oxygen is reduced. Barium and strontium behave similarly, giving BaO and SrO. Recognise that the oxide is ionic and basic, and that this is a redox reaction rather than simple combination alone.

This item also belongs to the Group 2 section and names water as the reagent. The task is to identify what a Group 2 metal forms when it reacts with water. Group 2 metals react with water to give the metal hydroxide and hydrogen gas, for example Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g). The metal is oxidised to a 2⁺ ion, hydrogen in water is reduced from 1⁺ to 0 in H₂, and the hydroxide ion carries a 1⁻ charge, so two hydroxide ions balance one 2⁺ metal ion. Calcium, strontium and barium react more vigorously than magnesium, and the resulting solutions are alkaline because the hydroxide ions are present. Write the balanced equation with state symbols and identify the gas evolved as hydrogen, which gives a squeaky pop with a lit splint. Distinguish this from the oxide formed with oxygen, and note that the hydroxide is basic.

This clause sits inside the Group 2 reactions list, so the assessed skill is predicting and describing how Group 2 metals react with dilute acids. A Group 2 metal such as magnesium reacts with a dilute acid such as hydrochloric acid to give a salt and hydrogen gas, for example Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). The metal is oxidised and the hydrogen ions in the acid are reduced, so the reaction is a redox process as well as a displacement of hydrogen. Reactivity increases down the group, so barium reacts more vigorously than calcium and calcium more vigorously than magnesium, matching the trend in first and second ionisation energies. Learners should balance equations, name the salt formed from the acid used, and identify the gas by its test.

This clause asks learners to explain why Group 2 metals become more reactive down the group using first and second ionisation energies. Going down the group, atomic radius increases and the outer electrons sit in a higher principal shell, so they are further from the nucleus and more shielded. Both the first and second ionisation energies therefore fall, meaning less energy is needed to remove the two outer electrons to form a 2⁺ ion. For example, magnesium has higher first and second ionisation energies than calcium, so calcium loses its two electrons more readily and reacts more vigorously. Learners should link the falling ionisation energies to the greater ease of forming the 2⁺ ion and hence the increased reactivity.

This row is a course and text-selection guide for the Group 2 topic, not an assessed answer target. Read it alongside the specification statements that set out the required reactions of Group 2 compounds, and use the wording of those statements to decide which reactions, observations and equations you must be able to recall and write. Build a summary table of each named compound, the reagent it reacts with, the observation and the balanced equation, including state symbols. Check each equation balances in atoms and charge, and practise writing ionic equations where a precipitate or neutralisation occurs. Use the specification's own command words to decide whether you must state, describe or explain, and revise only the reactions the specification lists rather than adding reactions from other sources.

Group 2 oxides are basic anhydrides: they react with water to form the metal hydroxide, for example CaO(s) + H₂O(l) → Ca(OH)₂(aq). The resulting solutions are alkaline because the hydroxide ion is released, so the approximate pH is above 7, typically around 10 to 12 for the more soluble hydroxides. Alkalinity increases down the group from magnesium to barium because the hydroxides become more soluble, so more hydroxide ions are present in solution. Magnesium oxide reacts only very slightly with water and magnesium hydroxide is only sparingly soluble, giving the weakest alkalinity; barium oxide reacts readily and barium hydroxide is the most soluble, giving the strongest alkalinity. When answering, write the equation, state that the solution is alkaline with a pH above 7, and explain the trend using increasing solubility of the hydroxide down the group.

Group 2 oxides and hydroxides act as bases because their oxide ions and hydroxide ions accept protons from acids. Calcium oxide (CaO) reacts with dilute hydrochloric acid to give calcium chloride and water: CaO + 2HCl → CaCl₂ + H₂O. Calcium hydroxide, Ca(OH)₂, neutralises acids in the same way: Ca(OH)₂ + 2HCl → CaCl₂ + 2H₂O. Magnesium hydroxide, Mg(OH)₂, works similarly and appears in indigestion remedies. The ionic view is that O²⁻ and OH⁻ accept H⁺ ions, so these compounds neutralise acids and raise pH. Write balanced equations with correct state symbols and check that atoms and charges balance on both sides.

Farmers spread calcium hydroxide, Ca(OH)₂, also called slaked lime, on acidic soil to raise its pH. The hydroxide ions accept protons from the acid in the soil, neutralising it and making conditions better for crops. A simplified equation is Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O, showing that each hydroxide ion takes a proton to form water. Calcium hydroxide is a base, so it reacts with the acid rather than adding to it. In multiple-choice questions, link the use to neutralisation, identify Ca(OH)₂ as the base, and check that the equation balances in atoms and charge.

Indigestion is caused by excess hydrochloric acid in the stomach. Both magnesium hydroxide, Mg(OH)₂, and calcium carbonate, CaCO₃, act as antacids because they are bases that neutralise this acid, raising the pH towards neutral and relieving discomfort. Mg(OH)₂ is a sparingly soluble hydroxide that neutralises acid: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. CaCO₃ is a carbonate that neutralises acid and releases carbon dioxide: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. The two equations differ in products, so a question may test which salt and which gas form. Recognise that the active ingredient is the base, not the tablet filler, and that the Group 2 ion (Mg²⁺ or Ca²⁺) is a spectator in the neutralisation.