Sulphur and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Sulphur and Its Inorganic Compounds

4 lessons

Lesson 4 of 4

Testing for Sulfur Compounds; Sulfur Oxides and the Environment

25 min

A volcano next door

Nyiragongo, one of the world's most active volcanoes, stands about 20 km north of Goma, the city that borders Rubavu on the shore of Lake Kivu. Its lava lake gives off sulfur dioxide day and night. Scientists at the Goma Volcano Observatory measure this gas to watch the volcano: in a study of 2020–2021 they found that when its sulfur dioxide output rose above about 9 kilotonnes (9000 tonnes) a day, the volcano was in a phase of intense activity. How do chemists detect sulfur compounds, and what harm do the sulfur oxides do once they are in the air?

Tests for sulfur compounds

In S2 you tested for the sulfate ion with acidified barium chloride. The same reagent, used step by step, also identifies sulfites, and sulfur dioxide has its own test (Lesson 2). Table 4.1 brings them together.

Table 4.1: Tests for sulfates, sulfites, sulfides and sulfur dioxide

To test for	Method	Positive result
sulfate, $\mathrm{SO_4^{2-}}$	add dilute hydrochloric acid, then barium chloride solution	white precipitate of barium sulfate that does not dissolve in the acid
sulfite, $\mathrm{SO_3^{2-}}$	add barium chloride solution, then dilute hydrochloric acid; test the gas	white precipitate that dissolves in the acid, giving sulfur dioxide that turns acidified potassium dichromate(VI) green
sulfite (solid)	warm with dilute hydrochloric acid	sulfur dioxide given off (dichromate turns green)
sulfide, $\mathrm{S^{2-}}$	add dilute hydrochloric acid	hydrogen sulfide, smelling of rotten eggs, which turns moist lead(II) ethanoate paper black
sulfur dioxide gas	acidified potassium dichromate(VI) paper or solution	orange → green (also turns acidified potassium manganate(VII) from purple to colourless)

The equations behind the tests:

$\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)} \qquad \mathrm{Ba^{2+}(aq) + SO_3^{2-}(aq) \rightarrow BaSO_3(s)}$

$\mathrm{BaSO_3(s) + 2HCl(aq) \rightarrow BaCl_2(aq) + H_2O(l) + SO_2(g)}$

$\mathrm{Pb^{2+}(aq) + H_2S(g) \rightarrow PbS(s) + 2H^+(aq)}$

Why add acid in the sulfate test? Barium chloride also gives white precipitates with sulfite and carbonate ions. Their barium salts dissolve in hydrochloric acid, but barium sulfate does not, so a precipitate that survives the acid can only be the sulfate (Figure 4.1).

Figure 4.1: Telling a sulfate, a sulfite and a carbonate apart with barium chloride.

Warning

Safety. Barium compounds and hydrogen sulfide are toxic. Hydrogen sulfide is dangerous even though you can smell it: at high concentration it deadens the sense of smell. Use small amounts, test for gases in a fume cupboard, and wash your hands after handling barium chloride.

Environmental issues of sulfur oxides

Sources. Sulfur dioxide enters the air naturally from volcanoes, and from human activity: burning fuels that contain sulfur (diesel, heavy fuel oil and coal) in engines, generators, factories and power stations, and roasting sulfide ores to extract metals. In the air, some of it is oxidised to sulfur trioxide.

Effects.

Acid rain. As you saw in S1, sulfur oxides dissolve in cloud water to form sulfurous and sulfuric acids: $\mathrm{SO_2(g) + H_2O(l) \rightarrow H_2SO_3(aq)}$ and $\mathrm{SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)}$ .
Buildings and metals. Acid rain dissolves limestone and marble, and the calcium sulfate formed crumbles away: $\mathrm{CaCO_3(s) + H_2SO_4(aq) \rightarrow CaSO_4(s) + H_2O(l) + CO_2(g)}$ . It speeds up the rusting of iron roofs, bridges and vehicles.
Plants and soils. Leaves turn yellow and die; acid washes nutrients such as calcium and magnesium out of the soil.
Lakes and rivers become more acidic, and fish and water insects die.
Health. Sulfur dioxide irritates the eyes and airways and can bring on asthma attacks and bronchitis, especially in children and older people.

Control.

Remove sulfur from fuels at the refinery, so that vehicles and generators burn low-sulfur fuel.
Remove sulfur dioxide from the waste gases of power stations and factories by flue-gas desulfurisation (Figure 4.2). The calcium sulfite formed is oxidised by air to calcium sulfate (gypsum), which is sold to make plasterboard: $\mathrm{2CaSO_3(s) + O_2(g) \rightarrow 2CaSO_4(s)}$ .
Capture sulfur dioxide at smelters and use it to make sulfuric acid, turning a pollutant into a product.
Add powdered limestone to acidified lakes and soils to neutralise the acid.

Figure 4.2: Removing sulfur dioxide from waste gases with limestone (flue-gas desulfurisation).

Definition

Flue-gas desulfurisation (FGD): removing sulfur dioxide from the waste gases (flue gases) of power stations and factories before they leave the chimney, usually by reaction with limestone or lime. Example: $\mathrm{CaCO_3(s) + SO_2(g) \rightarrow CaSO_3(s) + CO_2(g)}$ .

Worked example

Problem. Suppose Nyiragongo gives off 9000 tonnes of sulfur dioxide in one day. If all of it eventually became sulfuric acid in rain, what mass of sulfuric acid would that be? (S = 32, O = 16, H = 1)

Solution. Each $\mathrm{SO_2}$ molecule ( $M_{r} = 64$ ) becomes one $\mathrm{H_2SO_4}$ molecule ( $M_{r} = 98$ ).

$\text{mass of } \mathrm{H_2SO_4} = 9000\,\mathrm{t} \times \frac{98}{64} = 13\,800\,\mathrm{t}$

That is about 13 800 tonnes of acid from one day's gas, which shows why the air and rain near an active volcano can be strongly acidic.

Worked example

Problem. A factory burns 100 tonnes of coal containing 2.0% sulfur by mass. (a) What mass of sulfur dioxide is formed? (b) What mass of limestone (calcium carbonate) is needed to remove it all by flue-gas desulfurisation? (S = 32, O = 16, Ca = 40, C = 12)

Solution. Mass of sulfur $= 0.020 \times 100 = 2.0\,\mathrm{t}$ .

(a) $\mathrm{S(s) + O_2(g) \rightarrow SO_2(g)}$ : 32 g of S gives 64 g of $\mathrm{SO_2}$ , so the mass is $2.0 \times \frac{64}{32} = 4.0\,\mathrm{t}$ .

(b) $\mathrm{CaCO_3(s) + SO_2(g) \rightarrow CaSO_3(s) + CO_2(g)}$ : 64 g of $\mathrm{SO_2}$ needs 100 g of $\mathrm{CaCO_3}$ , so the mass is $4.0 \times \frac{100}{64} = 6.25\,\mathrm{t}$ , about 6.3 tonnes.

Notice that removing sulfur dioxide this way releases carbon dioxide: cleaner air locally, but a little more greenhouse gas.

Tip

Barium chloride, then acid. A white precipitate that stays in acid: sulfate. One that dissolves with a gas that turns dichromate green: sulfite. One that dissolves with a gas that only turns limewater milky: carbonate.

FuseSchool: Reducing Acid Rain Or Its Effects (5 min). Watch for the ways of cutting sulfur dioxide at its source and of repairing the damage, and match each to the list above.

✦

Practice

Describe how you would show that a solution contains sulfate ions. Write the ionic equation.
Why is dilute hydrochloric acid added in the sulfate test?
A white solid fizzes when warmed with dilute hydrochloric acid, and the gas turns acidified potassium dichromate(VI) green. Name the anion present and the gas.
How would you show that a black solid contains sulfide ions?
Name two human sources of sulfur dioxide and two harmful effects of it.
Write the equation for the reaction of limestone with sulfuric acid in acid rain, and explain why statues are damaged.
Evaluate flue-gas desulfurisation as a way of reducing pollution, giving one advantage and one disadvantage.
A generator burns 50 kg of diesel containing 0.30% sulfur by mass. What mass of sulfur dioxide does it release? (S = 32, O = 16)

Show answersHide answers

Add dilute hydrochloric acid, then barium chloride solution: a white precipitate that does not dissolve shows sulfate. $\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)}$ .
To remove carbonate and sulfite ions, which would also give white precipitates with barium ions; their barium salts dissolve in the acid, but barium sulfate does not.
Sulfite ion, $\mathrm{SO_3^{2-}}$ ; the gas is sulfur dioxide.
Add dilute hydrochloric acid: a gas smelling of rotten eggs (hydrogen sulfide) is given off, which turns moist lead(II) ethanoate paper black.
Sources (any two): burning diesel, heavy fuel oil or coal; roasting sulfide ores. Effects (any two): acid rain; damage to buildings and metals; harm to plants, lakes and fish; breathing problems and asthma.
$\mathrm{CaCO_3(s) + H_2SO_4(aq) \rightarrow CaSO_4(s) + H_2O(l) + CO_2(g)}$ . The calcium carbonate of limestone and marble is slowly converted to calcium sulfate, which crumbles and washes away, so fine details are lost.
Advantage (any one): removes most of the sulfur dioxide; the gypsum formed is a useful product. Disadvantage (any one): needs large amounts of limestone, which must be quarried; it releases carbon dioxide; it is expensive to build and run.
Sulfur $= 0.0030 \times 50 = 0.15\,\mathrm{kg}$ ; $\mathrm{SO_2} = 0.15 \times \frac{64}{32} = 0.30\,\mathrm{kg}$ (300 g).

1A volcano next door

2Tests for sulfur compounds

3Environmental issues of sulfur oxides

A volcano next door

Tests for sulfur compounds

Environmental issues of sulfur oxides