Oxoanions of Sulphur, Tests for Sulphite and Sulphate, and Uses of Group 16 – Trends in Chemical Properties of Group 16 Elements and Their Compounds (Senior 4 Chemistry) | Ganzaa

Trends in Chemical Properties of Group 16 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 16 Elements and Their Compounds

4 lessons

Lesson 4 of 4

Oxoanions of Sulphur, Tests for Sulphite and Sulphate, and Uses of Group 16

25 min

One element, many anions

Read the label on a bottle of fruit squash or a packet of dried fruit and you may see "preservative: sodium metabisulphite (E223)". In a photography darkroom, crystals of sodium thiosulphate "fix" the image; in hospitals, barium sulphate shows up the gut on X-rays. These are all salts of sulphur oxoanions, ions made of sulphur and oxygen, in which sulphur takes oxidation states from +2 to +6. This lesson describes the main oxoanions, how to tell sulphite from sulphate, and how Group 16 elements and compounds are used.

Definition

Oxoanion: a negative ion containing an element combined with oxygen. Example: sulphate, $\mathrm{SO_4^{2-}}$ , in which sulphur is +6.

Iodometric titration: a titration in which iodine produced (or used) in a reaction is measured with standard sodium thiosulphate solution: $\mathrm{2S_2O_3^{2-} + I_2 \rightarrow S_4O_6^{2-} + 2I^-}$ .

The main oxoanions

Table 11.3: Oxoanions of sulphur

Ion	Formula	Oxidation state of S	Typical behaviour
sulphite	$\mathrm{SO_3^{2-}}$	+4	reducing agent; gives $\mathrm{SO_2}$ with acid
hydrogensulphite	$\mathrm{HSO_3^-}$	+4	as sulphite; formed when $\mathrm{SO_2}$ dissolves in water
disulphite (metabisulphite)	$\mathrm{S_2O_5^{2-}}$	+4	releases $\mathrm{SO_2}$ : food and drink preservative
sulphate	$\mathrm{SO_4^{2-}}$	+6	very stable; most metal sulphates soluble, but $\mathrm{BaSO_4}$ insoluble
hydrogensulphate	$\mathrm{HSO_4^-}$	+6	acidic in water
thiosulphate	$\mathrm{S_2O_3^{2-}}$	average +2	decomposes in acid to S and $\mathrm{SO_2}$ ; reduces iodine
peroxodisulphate	$\mathrm{S_2O_8^{2-}}$	+6 (contains an O–O link)	very strong oxidising agent

Figure 11.7: Sulphite, sulphate and thiosulphate.

Key reactions

Sulphite as a reducing agent. Sulphite is oxidised to sulphate, $\mathrm{SO_3^{2-} + H_2O \rightarrow SO_4^{2-} + 2H^+ + 2e^-}$ , so it decolourises acidified potassium manganate(VII) and turns acidified dichromate(VI) from orange to green. Sulphite solutions slowly turn into sulphate in air, which is why sulphite preservatives are used up over time.
Sulphite with acid gives sulphur dioxide: $\mathrm{SO_3^{2-} + 2H^+ \rightarrow SO_2 + H_2O}$ .
Thiosulphate with acid slowly forms a pale yellow cloud of sulphur, $\mathrm{S_2O_3^{2-} + 2H^+ \rightarrow S + SO_2 + H_2O}$ : the classic experiment for studying reaction rates.
Thiosulphate with iodine: $\mathrm{2S_2O_3^{2-} + I_2 \rightarrow S_4O_6^{2-} + 2I^-}$ (tetrathionate). The brown iodine colour disappears, and starch gives a sharp end point: the basis of iodometric titrations.
Peroxodisulphate oxidises iodide to iodine: $\mathrm{S_2O_8^{2-} + 2I^- \rightarrow 2SO_4^{2-} + I_2}$ .

Tests for sulphate and sulphite ions

Sulphate. Acidify the solution with dilute hydrochloric acid, then add barium chloride solution. A white precipitate of barium sulphate that does not dissolve in the acid shows sulphate: $\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)}$ . The acid is added first to remove carbonate and sulphite ions, which would also give white barium salts.
Sulphite. Barium chloride gives a white precipitate of barium sulphite, but it dissolves in dilute hydrochloric acid, giving sulphur dioxide: $\mathrm{BaSO_3(s) + 2H^+(aq) \rightarrow Ba^{2+}(aq) + SO_2(g) + H_2O(l)}$ . Sulphur dioxide has a choking smell and turns filter paper soaked in acidified potassium dichromate(VI) from orange to green (or decolourises acidified manganate(VII)).

Figure 11.8: Telling sulphate, sulphite and carbonate apart.

Colourless crystals of sodium thiosulphate (as the pentahydrate, Na₂S₂O₃·5H₂O) in a small bottle. Solutions of it are used in iodometric titrations and, in film photography, to dissolve away unexposed silver halide. Photo: W. Oelen, CC BY-SA 3.0, via Wikimedia Commons

Worked example 1: measuring sulphate by mass

Excess acidified barium chloride is added to 25.0 cm³ of a solution containing sulphate ions. The precipitate is filtered, washed, dried and weighed: 0.233 g. Calculate the concentration of sulphate ions. $M(\mathrm{BaSO_4}) = 233.4\ \mathrm{g\,mol^{-1}}$ .

Solution. 1 mol $\mathrm{SO_4^{2-}}$ gives 1 mol $\mathrm{BaSO_4}$ . $n = \frac{0.233}{233.4} = 9.98 \times 10^{-4}\ \mathrm{mol} \qquad c = \frac{9.98 \times 10^{-4}}{0.0250} = 0.0399\ \mathrm{mol\,dm^{-3}}$ Barium sulphate is so insoluble that almost all the sulphate is collected, which makes this an accurate method (gravimetric analysis).

Worked example 2: an iodometric titration

25.0 cm³ of an iodine solution needs 20.0 cm³ of 0.100 mol dm⁻³ sodium thiosulphate to remove the iodine colour (starch indicator). Calculate the concentration of iodine.

Solution. $\mathrm{2S_2O_3^{2-} + I_2 \rightarrow S_4O_6^{2-} + 2I^-}$ . $n(\mathrm{S_2O_3^{2-}}) = 0.0200 \times 0.100 = 2.00 \times 10^{-3}\ \mathrm{mol} \qquad n(\mathrm{I_2}) = \frac{2.00 \times 10^{-3}}{2} = 1.00 \times 10^{-3}\ \mathrm{mol}$ $c(\mathrm{I_2}) = \frac{1.00 \times 10^{-3}}{0.0250} = 0.0400\ \mathrm{mol\,dm^{-3}}$

Uses of Group 16 elements and compounds

Table 11.4: Uses and the properties behind them

Element or compound	Use	Property that makes it suitable
Oxygen	medical oxygen; steelmaking (blowing impurities out of iron); welding with ethyne	supports respiration and burning; oxidises carbon and other impurities
Ozone	disinfecting drinking water	very strong oxidising agent that kills microbes
Sulphur	making sulphuric acid; vulcanising rubber; fungicide on crops	burns to $\mathrm{SO_2}$ ; forms S–S cross-links between rubber chains; toxic to fungi
Sulphuric acid	fertilisers (superphosphate, ammonium sulphate); car batteries; detergents	strong, involatile acid; electrolyte
Sulphur dioxide and sulphites	preserving dried fruit, juices and wine; bleaching paper pulp	reducing agents that kill microbes and remove coloured compounds
Sodium thiosulphate	iodometric titrations; photographic fixer	reduces iodine; dissolves silver halides
Barium sulphate	"barium meal" for X-rays	insoluble (so not poisonous) and absorbs X-rays
Selenium and tellurium	solar cells (cadmium telluride); photocells; colouring and decolourising glass	semiconductors whose conductivity changes with light

A hazard: acid rain. Burning fuels that contain sulphur releases sulphur dioxide, which is oxidised in the air and dissolves in rain as sulphuric acid. Acid rain damages forests, lakes and buildings, which is why sulphur is removed from fuels and from power-station exhaust gases.

Sodium metabisulphite (sodium disulphite, Na₂S₂O₅), a white powder. It releases sulphur dioxide in water and acid, which is why it is used as a preservative (E223) in some drinks and dried fruits. Photo: Walkerma, public domain, via Wikimedia Commons

Warning

Common mistakes. (1) Adding barium chloride without acidifying first: carbonate and sulphite also give white precipitates. (2) Using dilute sulphuric acid to acidify a sulphate test: it adds sulphate ions itself. (3) Identifying $\mathrm{SO_2}$ with limewater alone: carbon dioxide also turns it milky; use acidified dichromate. (4) Writing thiosulphate as $\mathrm{SO_3^{2-}}$ or $\mathrm{S_2O_4^{2-}}$ : it is $\mathrm{S_2O_3^{2-}}$ .

Safety. Soluble barium salts such as barium chloride are poisonous if swallowed: use small volumes and wash hands afterwards. Sulphur dioxide is a toxic, choking gas that can trigger asthma attacks: generate it only in small amounts in a fume cupboard. Potassium dichromate is toxic and a carcinogen, so only test papers or very dilute solutions should be used, with gloves. Wear eye protection throughout.

Tip

Sulphate stays, sulphite goes. With $\mathrm{BaCl_2}$ then HCl: $\mathrm{BaSO_4}$ stays; $\mathrm{BaSO_3}$ dissolves and gives $\mathrm{SO_2}$ (orange dichromate → green). Thiosulphate + acid = yellow sulphur; thiosulphate + iodine = colourless.

Digital Science Uganda: Test for sulphite and sulphate ions (4 min). Watch how barium chloride and dilute acid are used to tell the two ions apart, and how the gas from the sulphite is identified.

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Practice

Give the oxidation state of sulphur in $\mathrm{SO_3^{2-}}$ , $\mathrm{SO_4^{2-}}$ , $\mathrm{S_2O_3^{2-}}$ and $\mathrm{S_4O_6^{2-}}$ .
Describe how you would show that a solution contains sulphate ions. Why is the acid added?
Describe how you would distinguish a solution of sodium sulphite from one of sodium sulphate.
Write equations for: (a) sulphite ions reacting with dilute acid; (b) thiosulphate ions reacting with dilute acid; (c) thiosulphate ions reacting with iodine.
Why can limewater not be used on its own to identify sulphur dioxide?
For each use, name the Group 16 substance and the property behind it: (a) X-ray "meals"; (b) dried-fruit preservative; (c) vulcanising rubber; (d) disinfecting water without chlorine.
0.466 g of barium sulphate is obtained from 50.0 cm³ of a water sample. Calculate the sulphate concentration in mol dm⁻³. $M(\mathrm{BaSO_4}) = 233.4$ .
In $\mathrm{2S_2O_3^{2-} + I_2 \rightarrow S_4O_6^{2-} + 2I^-}$ , show that the electrons lost equal the electrons gained.

Show answersHide answers

+4; +6; average +2; average +2.5.
Add dilute hydrochloric (or nitric) acid, then barium chloride (or nitrate) solution: a white precipitate that does not dissolve shows sulphate. The acid removes carbonate and sulphite ions, which would also give white precipitates.
Add barium chloride to each, then dilute HCl: the sulphate precipitate stays; the sulphite precipitate dissolves with fizzing, and the gas ( $\mathrm{SO_2}$ ) turns acidified dichromate paper from orange to green. (Or: sulphite decolourises acidified manganate(VII); sulphate does not.)
(a) $\mathrm{SO_3^{2-} + 2H^+ \rightarrow SO_2 + H_2O}$ ; (b) $\mathrm{S_2O_3^{2-} + 2H^+ \rightarrow S + SO_2 + H_2O}$ ; (c) $\mathrm{2S_2O_3^{2-} + I_2 \rightarrow S_4O_6^{2-} + 2I^-}$ .
Carbon dioxide also turns limewater milky; acidified dichromate (orange → green) responds to $\mathrm{SO_2}$ but not $\mathrm{CO_2}$ .
(a) barium sulphate: insoluble and absorbs X-rays; (b) sulphur dioxide or a sulphite/metabisulphite: kills microbes, reducing agent; (c) sulphur: forms S–S cross-links between rubber chains; (d) ozone: strong oxidising agent that kills microbes.
$n = \frac{0.466}{233.4} = 2.00 \times 10^{-3}$ mol; $c = \frac{2.00 \times 10^{-3}}{0.0500} = 0.0400\ \mathrm{mol\,dm^{-3}}$ .
Iodine: $\mathrm{I_2 + 2e^- \rightarrow 2I^-}$ , gains 2 electrons. Sulphur: average +2 in $\mathrm{S_2O_3^{2-}}$ to +2.5 in $\mathrm{S_4O_6^{2-}}$ ; four S atoms each rise by 0.5, a loss of 2 electrons. 2 lost = 2 gained.

Quick check

A white precipitate forms when barium chloride is added to a solution. It dissolves in dilute HCl with a gas that turns acidified dichromate paper green. Which ion was present?

Key points

Sulphur oxoanions: sulphite $\mathrm{SO_3^{2-}}$ (+4, reducing), sulphate $\mathrm{SO_4^{2-}}$ (+6, stable), thiosulphate $\mathrm{S_2O_3^{2-}}$ (gives S with acid, reduces $\mathrm{I_2}$ ), peroxodisulphate $\mathrm{S_2O_8^{2-}}$ (strong oxidant).
Sulphate test: dilute HCl then $\mathrm{BaCl_2}$ : white $\mathrm{BaSO_4}$ , insoluble in acid.
Sulphite test: white $\mathrm{BaSO_3}$ dissolves in dilute HCl giving $\mathrm{SO_2}$ , which turns acidified dichromate from orange to green.
Uses follow properties: $\mathrm{O_2}$ (medicine, steel), $\mathrm{O_3}$ (water), S ( $\mathrm{H_2SO_4}$ , rubber), $\mathrm{SO_2}$ /sulphites (preservatives, bleach), thiosulphate (titrations), $\mathrm{BaSO_4}$ (X-rays), Se/Te (solar cells). $\mathrm{SO_2}$ causes acid rain.

1One element, many anions

2The main oxoanions

3Key reactions

4Tests for sulphate and sulphite ions

5Uses of Group 16 elements and compounds

One element, many anions

The main oxoanions

Key reactions

Tests for sulphate and sulphite ions

Uses of Group 16 elements and compounds