The alkaline earth metals: magnesium and calcium: WASSCE Chemistry

Magnesium burns with a dazzling white light and lightens the alloys in aircraft; calcium builds bones, teeth and seashells. They belong to group 2, the alkaline earth metals. This lesson covers their properties and trends, the reactions of magnesium and calcium, and how they compare with the alkali metals of group 1.

Learning objectives

By the end of this lesson you should be able to:

name the alkaline earth metals and relate their properties to their electron configuration
compare group 2 metals with group 1 metals
describe and write equations for the reactions of magnesium and calcium with oxygen, water, steam and acids
state and explain the trend in reactivity down group 2
describe the extraction and uses of magnesium and calcium
state the biological importance of magnesium and calcium

The alkaline earth metals

Group 2: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra).

Element	Configuration	Ion
Beryllium	2, 2	Be²⁺
Magnesium	2, 8, 2	Mg²⁺
Calcium	2, 8, 8, 2	Ca²⁺

Each has two outer electrons, which it loses to form a 2+ ion. Their oxides and hydroxides are basic. The name "alkaline earth" comes from the fact that their oxides are alkaline and were once called "earths".

Group 2 compared with group 1

Property	Group 1 (e.g. Na)	Group 2 (e.g. Mg)
Outer electrons	1	2
Ion	M⁺	M²⁺
Hardness	Soft (cut with a knife)	Harder
Melting point	Low	Higher (stronger metallic bonding: two electrons per atom)
Density	Lower	Higher
Reactivity	Very reactive	Less reactive (two electrons must be removed)
Reaction with cold water	Vigorous	Mg very slow; Ca steady
Carbonates on heating	Stable	Decompose to oxide + CO₂
Hydroxides	Very soluble, strong alkalis	Less soluble (Ca(OH)₂ slightly soluble; Mg(OH)₂ insoluble)

Reactions of magnesium and calcium

With oxygen

Magnesium burns with a dazzling white flame, leaving white magnesium oxide:

2\text{Mg(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MgO(s)}

\begin{gathered}2\text{Mg(s)} + \text{O}_2\text{(g)} \\[6pt] \rightarrow 2\text{MgO(s)}\end{gathered}

Calcium burns with a brick-red flame: 2Ca + O₂ → 2CaO.

With water and steam

Metal	Cold water	Steam
Magnesium	Very slow (bubbles over days): Mg + 2H₂O → Mg(OH)₂ + H₂	Burns brightly: Mg + H₂O(g) → MgO + H₂
Calcium	Steady fizzing; sinks; solution turns cloudy (Ca(OH)₂ is slightly soluble): Ca + 2H₂O → Ca(OH)₂ + H₂	Vigorous

With dilute acids

Both react rapidly, giving a salt and hydrogen:

\text{Mg(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + \text{H}_2\text{(g)}

\begin{gathered}\text{Mg(s)} + 2\text{HCl(aq)} \\[6pt] \rightarrow \text{MgCl}_2\text{(aq)} + \text{H}_2\text{(g)}\end{gathered}

Calcium with dilute sulphuric acid soon stops, because insoluble calcium sulphate coats the metal.

With carbon(IV) oxide and nitrogen

Burning magnesium continues to burn in carbon(IV) oxide, leaving black specks of carbon: 2Mg + CO₂ → 2MgO + C. It also combines with nitrogen when burned in air, forming some magnesium nitride: 3Mg + N₂ → Mg₃N₂.

Trend in reactivity

Reactivity increases down the group (Be < Mg < Ca < Sr < Ba). The two outer electrons are further from the nucleus and more shielded in larger atoms, so they are lost more easily. Barium reacts with cold water about as vigorously as sodium.

Extraction

Both metals are very reactive, so they are extracted by the electrolysis of their molten chlorides:

Magnesium: from sea water. Magnesium ions are precipitated as Mg(OH)₂ (using lime), converted to MgCl₂ with HCl, dried and electrolysed: Mg²⁺ + 2e⁻ → Mg (cathode); 2Cl⁻ → Cl₂ + 2e⁻ (anode).
Calcium: electrolysis of molten calcium chloride.

Uses

Magnesium	Calcium
Light alloys (with aluminium) for aircraft, car wheels and bicycles	Removing oxygen and sulphur from steel
Flares, fireworks and tracer bullets (bright white light)	Extracting other metals (e.g. uranium)
Sacrificial anodes to protect ships and pipelines from rusting	Its compounds: lime, cement, plaster (next lesson)
Magnesium hydroxide ("milk of magnesia") as an antacid and laxative

In living things

Calcium: bones and teeth (calcium phosphate), shells (calcium carbonate); needed for blood clotting and muscle contraction. Good sources: milk, small fish eaten whole, leafy greens.
Magnesium: at the centre of the chlorophyll molecule, so plants cannot photosynthesise without it; also needed by enzymes in the body.

In Sierra Leone

Small dried fish such as bonga and the tiny fish eaten whole in Sierra Leonean sauces are an excellent source of calcium, because their soft bones are eaten too. Palm-tree and other crops showing yellowing between leaf veins may be short of magnesium, because magnesium is part of chlorophyll. Agricultural lime containing magnesium (dolomite) can correct it.

Common mistake

Group 2 metals are less reactive than group 1 metals in the same period: removing two electrons needs more energy than removing one. Magnesium is much less reactive than sodium.

Key points

Group 2 metals have two outer electrons and form M²⁺ ions.
They are harder, denser, higher-melting and less reactive than group 1 metals.
Mg burns with a dazzling white flame, reacts very slowly with cold water but burns in steam; Ca reacts steadily with cold water.
Reactivity increases down the group.
Extracted by electrolysis of molten chlorides; Mg comes from sea water.
Uses: Mg alloys, flares, sacrificial anodes, antacids. Ca is vital for bones and teeth; Mg is in chlorophyll.

WASSCE examination focus

WAEC tests this

Expect to compare group 1 and group 2 metals, to describe the reactions of magnesium and calcium with water, steam, oxygen and acids (with equations), to explain the trend in reactivity, and to state uses of magnesium and its compounds.

Quick check

Why are group 2 metals less reactive than group 1 metals?
Write the equation for magnesium reacting with steam.
Why does the reaction of calcium with dilute sulphuric acid soon stop?
Give two uses of magnesium.

Answers:

They must lose two electrons, which needs more energy than losing one.
Mg + H₂O(g) → MgO + H₂.
Insoluble calcium sulphate forms a coating on the calcium.
Any two: light alloys, flares and fireworks, sacrificial anodes, antacids (as Mg(OH)₂).

The alkaline earth metals: magnesium and calcium