Trends in Chemical Properties of Group 2 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 2 Elements and Their Compounds

5 lessons

Lesson 5 of 5

Uses of Group 2 Elements and Compounds

25 min

Sweetening sour soil

Acidic soils cover about 45% of Rwanda's arable land and about 60% of its highland farmland. Below a pH of about 5.5, crops struggle: phosphorus becomes less available and aluminium and manganese ions reach toxic levels around the roots. The remedy is a Group 2 compound. Farmers spread crushed travertine, an uncalcined form of limestone (calcium carbonate), and the Rwanda Agriculture Board recommends 2.5 tonnes per hectare once every two years. In this lesson you will see how the properties you have studied in this unit, basic oxides, insoluble sulphates, light and reactive metals, explain the many uses of Group 2 elements and their compounds.

Definition

Agricultural lime: a calcium compound, usually crushed limestone or travertine ( $\mathrm{CaCO_3}$ ), sometimes slaked lime ( $\mathrm{Ca(OH)_2}$ ) or quicklime (CaO), spread on soil to neutralise acidity and raise its pH.

Sacrificial protection: protecting iron or steel from rusting by connecting it to a more reactive metal, such as magnesium, which corrodes instead of the iron.

Refractory: a material that keeps its strength at very high temperatures, used to line furnaces and kilns. Example: magnesium oxide, melting point 2852 °C.

Neutralising acids: limestone and its products

Calcium carbonate, calcium oxide and calcium hydroxide are all bases. In soil they remove hydrogen ions: $\mathrm{CaCO_3(s) + 2H^{+}(aq) \rightarrow Ca^{2+}(aq) + H_2O(l) + CO_2(g)}$ $\mathrm{Ca(OH)_2(s) + 2H^{+}(aq) \rightarrow Ca^{2+}(aq) + 2H_2O(l)}$ As the pH rises, aluminium ions are precipitated as aluminium hydroxide and stop harming roots, and the calcium is itself a plant nutrient. A study of travertine on Rwandan maize and bean fields found that liming paid off mainly where the soil pH was below about 5.2, and mostly when fertiliser was used as well: testing soil before liming saves money.

The same three compounds are linked by the lime cycle (Figure 7.10), which you can follow with the reactions from Lessons 2 and 3. Lime mortar, used for centuries in building, is slaked lime mixed with sand and water; it hardens slowly as it absorbs carbon dioxide from the air and turns back into calcium carbonate.

Figure 7.10: The lime cycle. Drawn by Ganzaa.

Cement in Rwanda

Cement is made by heating limestone with clay in a kiln at about 1450 °C. The calcium carbonate decomposes, and the calcium oxide combines with silica and alumina from the clay to form calcium silicates and aluminates, ground into cement. CIMERWA, at Bugarama in Rusizi District, is Rwanda's only integrated cement producer: it mines its own raw materials, makes the clinker and sells the cement. Every tonne of limestone decomposed releases about 440 kg of carbon dioxide (Worked example 2), which is why cement making is a major source of this greenhouse gas worldwide.

Uses across the group

Table 7.8: Uses of Group 2 elements and compounds, and the properties behind them

Substance	Use	Property it depends on
Magnesium alloys	light parts for aircraft, cars, bicycles and laptop cases	low density (1.74 g cm⁻³); strong when alloyed with aluminium and zinc
Magnesium metal	sacrificial anodes on steel pipelines, ship hulls and water heaters	more reactive than iron, so it is oxidised instead of the iron
Magnesium powder	flares and the white sparks in fireworks	burns with an intense white light
Magnesium oxide	lining furnaces and kilns	refractory: melting point 2852 °C
Magnesium hydroxide	antacid ("milk of magnesia")	a weak base, almost insoluble, so it neutralises stomach acid gently
Calcium carbonate	cement, building stone, agricultural lime, flux in the iron blast furnace	decomposes to CaO on heating; neutralises acids
Calcium oxide and hydroxide	mortar, whitewash, liming soil, treating water and acidic waste	cheap, strong bases
Calcium sulphate (plaster of Paris, $\mathrm{CaSO_4{\cdot}\frac{1}{2}H_2O}$ )	plaster for walls and casts for broken bones	takes up water again and sets hard as gypsum
Strontium carbonate and nitrate	red fireworks and red distress flares	crimson flame colour
Barium sulphate	"barium meal" for X-rays of the gut; white pigment; dense drilling mud	almost insoluble so not toxic; barium absorbs X-rays; high density
Beryllium and its alloys	telescope mirrors; non-sparking copper–beryllium tools; windows in X-ray tubes	light and stiff; keeps its shape when cold; transparent to X-rays

Worked example 1: an antacid tablet

A tablet contains 0.300 g of magnesium hydroxide. What volume of hydrochloric acid of concentration 0.100 mol dm⁻³ (similar to stomach acid) can it neutralise? $M(\mathrm{Mg(OH)_2}) = 58.32\ \mathrm{g\,mol^{-1}}$ .

Solution. $\mathrm{Mg(OH)_2 + 2HCl \rightarrow MgCl_2 + 2H_2O}$ $n(\mathrm{Mg(OH)_2}) = \frac{0.300}{58.32} = 5.14 \times 10^{-3}\ \mathrm{mol} \qquad n(\mathrm{HCl}) = 2 \times 5.14 \times 10^{-3} = 1.03 \times 10^{-2}\ \mathrm{mol}$ $V = \frac{1.03 \times 10^{-2}\ \mathrm{mol}}{0.100\ \mathrm{mol\,dm^{-3}}} = 0.103\ \mathrm{dm^3} = 103\ \mathrm{cm^3}$

Worked example 2: carbon dioxide from a lime kiln

Calculate the mass of carbon dioxide released when 1.00 tonne of pure calcium carbonate is decomposed. $M(\mathrm{CaCO_3}) = 100.09$ , $M(\mathrm{CO_2}) = 44.01\ \mathrm{g\,mol^{-1}}$ .

Solution. $\mathrm{CaCO_3 \rightarrow CaO + CO_2}$ : 100.09 g of carbonate gives 44.01 g of carbon dioxide, so the mass ratio is fixed. $m(\mathrm{CO_2}) = 1.00\ \mathrm{t} \times \frac{44.01}{100.09} = 0.440\ \mathrm{t} = 440\ \mathrm{kg}$

Worked example 3: evaluating a liming material

A cooperative can buy crushed limestone ( $\mathrm{CaCO_3}$ ) or quicklime (CaO). Compare the amount of acid each kilogram can neutralise, and evaluate which is the better choice for smallholders.

Solution. Each mole of either compound neutralises 2 mol of $\mathrm{H^+}$ . $\mathrm{CaCO_3}: \ \frac{1000}{100.09} \times 2 = 20.0\ \mathrm{mol\,H^+\,kg^{-1}} \qquad \mathrm{CaO}: \ \frac{1000}{56.08} \times 2 = 35.7\ \mathrm{mol\,H^+\,kg^{-1}}$ Quicklime neutralises about 1.8 times as much acid per kilogram, so less has to be carried. But it is caustic, burns skin and eyes, heats up violently with rain, and must itself be made by heating limestone, which costs fuel and releases carbon dioxide. Crushed limestone or travertine acts more slowly but is safe to handle, cheap where it is quarried nearby, and cannot "burn" a crop. For smallholders spreading lime by hand, limestone is the better choice; quicklime suits industry, where it can be handled safely.

Warning

Common mistakes. (1) Saying barium meals are safe because "barium is not toxic": soluble barium compounds are poisonous; only the insolubility of $\mathrm{BaSO_4}$ makes it safe. (2) Saying magnesium protects iron by "coating" it: sacrificial protection works because magnesium is more reactive and is oxidised first. (3) Confusing quicklime (CaO) with slaked lime ( $\mathrm{Ca(OH)_2}$ ).

Safety. Quicklime and slaked lime are corrosive to eyes and skin: wear goggles and gloves when spreading lime, and never rub lime dust from your eyes; rinse them with plenty of clean water and get medical help.

Tip

Link every use to a property: light (Mg, Be alloys), reactive (Mg anodes), basic (CaCO₃, CaO, Ca(OH)₂, Mg(OH)₂), insoluble (BaSO₄), coloured flame (Sr). Exam answers that name the property score; a bare list of uses does not.

FuseSchool: Limestone Cycle, limestone, quicklime and slaked lime (5 min). Follow each step of the cycle in Figure 7.10 and note the conditions and the observations at each step.

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Practice

Write an equation for the reaction of magnesium hydroxide with hydrochloric acid in the stomach.
Explain why magnesium blocks are attached to the steel hulls of ships.
Give the property of barium sulphate that makes it safe for a barium meal, and the property that makes it useful for X-rays.
Write equations for the three steps that turn limestone into limewater, and for the reaction that turns limewater milky.
Plaster of Paris is $\mathrm{CaSO_4{\cdot}\frac{1}{2}H_2O}$ . Write an equation for its setting with water to form gypsum, $\mathrm{CaSO_4{\cdot}2H_2O}$ .
A field of 0.40 hectares is limed at 2.5 tonnes per hectare with travertine. Assuming the travertine is pure calcium carbonate, calculate the amount, in moles, of $\mathrm{H^+}$ it can neutralise.
Evaluate the use of magnesium rather than aluminium for car wheels, using data from this unit (density of aluminium 2.70 g cm⁻³).

Show answersHide answers

$\mathrm{Mg(OH)_2(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + 2H_2O(l)}$
Magnesium is more reactive than iron, so it is oxidised in sea water instead of the steel (sacrificial protection); the blocks are replaced when they are used up.
Safe: almost insoluble in water and in stomach acid, so very few toxic $\mathrm{Ba^{2+}}$ ions are absorbed. Useful: barium is a heavy element that absorbs X-rays strongly, so the gut shows up white.
$\mathrm{CaCO_3 \rightarrow CaO + CO_2}$ ; $\mathrm{CaO + H_2O \rightarrow Ca(OH)_2}$ ; $\mathrm{Ca(OH)_2(s) \rightarrow Ca(OH)_2(aq)}$ (dissolving in more water, then filtering); $\mathrm{Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)}$ .
$\mathrm{CaSO_4{\cdot}\frac{1}{2}H_2O(s) + \frac{3}{2}H_2O(l) \rightarrow CaSO_4{\cdot}2H_2O(s)}$ , or doubled: $\mathrm{2CaSO_4{\cdot}\frac{1}{2}H_2O + 3H_2O \rightarrow 2CaSO_4{\cdot}2H_2O}$ .
Mass $= 0.40 \times 2.5 = 1.0$ t $= 1.0 \times 10^6$ g; $n(\mathrm{CaCO_3}) = \frac{1.0 \times 10^6}{100.09} = 9.99 \times 10^3$ mol; $n(\mathrm{H^+}) = 2.0 \times 10^4$ mol.
Magnesium (1.74 g cm⁻³) is about two thirds the density of aluminium (2.70 g cm⁻³), so a wheel of the same volume is lighter, saving fuel. But magnesium is more reactive: it corrodes more easily and burns if it catches fire, so it must be alloyed and coated, which costs more. Magnesium suits racing and luxury cars; aluminium is the cheaper everyday choice.

Quick check

Why is barium sulphate, but not barium chloride, safe to swallow before an X-ray?

Key points

Limestone, quicklime and slaked lime are bases used to lime acidic soils (about 45% of Rwanda's arable land), make cement and mortar, and treat water; they are linked by the lime cycle.
Magnesium: light alloys, sacrificial anodes, flares; MgO is refractory; $\mathrm{Mg(OH)_2}$ is an antacid.
Plaster of Paris sets by taking up water; strontium salts colour fireworks red; insoluble $\mathrm{BaSO_4}$ is used for X-ray meals.
Beryllium is light, stiff and transparent to X-rays: telescope mirrors, X-ray windows, non-sparking alloys.

1Sweetening sour soil

2Neutralising acids: limestone and its products

3Cement in Rwanda

4Uses across the group

Sweetening sour soil

Neutralising acids: limestone and its products

Cement in Rwanda

Uses across the group