Hydrides of Group 16: Acidity and Volatility – 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

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Senior 4 · Chemistry

Trends in Chemical Properties of Group 16 Elements and Their Compounds

4 lessons

Lesson 2 of 4

Hydrides of Group 16: Acidity and Volatility

25 min

Water and its poisonous cousin

Water, $\mathrm{H_2O}$ , is a liquid that boils at 100 °C, has a neutral pH and is essential to life. Hydrogen sulphide, $\mathrm{H_2S}$ , its neighbour one place down, is a gas that smells of rotten eggs, is weakly acidic in water, and is highly poisonous. Both molecules have the same bent shape with two hydrogens on a Group 16 atom, so why are they so different? This lesson compares the four hydrides from water to hydrogen telluride and explains two trends: how easily they boil (volatility) and how strongly they act as acids.

Definition

Volatility: how readily a substance turns into vapour. A more volatile substance has a lower boiling point. Example: $\mathrm{H_2S}$ (boiling point −59.6 °C) is far more volatile than water (100 °C).

Acid dissociation constant, $K_{a}$ , and $\mathrm{p}K_a = -\log_{10} K_a$ : measures of how far an acid gives up its proton in water. A smaller $\mathrm{p}K_a$ means a stronger acid. Example: $\mathrm{H_2S \rightleftharpoons H^+ + HS^-}$ , $\mathrm{p}K_{a1} = 7.0$ .

The data

Table 11.2: Hydrides of the Group 16 elements

Hydride	Boiling point / °C	$\mathrm{p}K_{a1}$ in water	Thermal stability	As a reducing agent
water, $\mathrm{H_2O}$	100	neutral (pH 7)	very stable	not reducing
hydrogen sulphide, $\mathrm{H_2S}$	−59.6	7.0	decomposes on strong heating	reducing
hydrogen selenide, $\mathrm{H_2Se}$	−41.3	3.89	less stable	more reducing
hydrogen telluride, $\mathrm{H_2Te}$	−2.2	2.6	decomposes easily	strongly reducing

Volatility: London forces and one big exception

From $\mathrm{H_2S}$ to $\mathrm{H_2Te}$ the boiling point rises (Figure 11.4). The molecules contain more electrons, so the temporary dipoles are larger and the London forces between molecules are stronger; the hydrides become less volatile down the group.

Water should, on this trend, boil far below −60 °C. Instead it boils at 100 °C, because oxygen is small and very electronegative and has two lone pairs: each water molecule can form up to four hydrogen bonds (two through its hydrogens, two through its lone pairs). Sulphur is much less electronegative (2.58 against 3.44), so $\mathrm{H_2S}$ forms no significant hydrogen bonds. The same pattern appears in Group 15 ( $\mathrm{NH_3}$ ) and Group 17 (HF).

Figure 11.4: Boiling points of the Group 16 hydrides.

Acidity: weaker bonds win

The hydrides become stronger acids down the group (Figure 11.5), even though the electronegativity of the central atom falls. The deciding factor is the H–X bond: as the atom X gets larger, the bond becomes longer and weaker, so the proton is lost more easily; the larger anion ( $\mathrm{HTe^-}$ compared with $\mathrm{HS^-}$ ) also spreads its negative charge over a bigger volume, so it holds the proton less tightly.

\mathrm{H_2S(aq) \rightleftharpoons H^+(aq) + HS^-(aq)} \qquad \mathrm{HS^-(aq) \rightleftharpoons H^+(aq) + S^{2-}(aq)}

Figure 11.5: First $\mathrm{p}K_a$ of the Group 16 hydrides; lower bars mean stronger acids.

The same weakening of the H–X bond makes the heavier hydrides less thermally stable and better reducing agents: hydrogen telluride decomposes to its elements easily, and hydrogen sulphide reduces chlorine, sulphur dioxide and acidified manganate(VII).

Hydrogen sulphide in the laboratory

Preparation. Dilute hydrochloric acid on iron(II) sulphide, traditionally in a Kipp's apparatus, which delivers gas only while the tap is open:

\mathrm{FeS(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2S(g)}

Reactions.

Burning: in excess air to sulphur dioxide, $\mathrm{2H_2S + 3O_2 \rightarrow 2SO_2 + 2H_2O}$ ; in limited air to sulphur, $\mathrm{2H_2S + O_2 \rightarrow 2S + 2H_2O}$ .
As a reducing agent it is oxidised to yellow sulphur: $\mathrm{H_2S + Cl_2 \rightarrow 2HCl + S}$ ; $\mathrm{2H_2S + SO_2 \rightarrow 3S + 2H_2O}$ (the basis of sulphur recovery from natural gas); it decolourises acidified potassium manganate(VII) and turns acidified dichromate(VI) from orange to green.
Precipitating metal sulphides: many metal ions give coloured, insoluble sulphides, for example black lead(II) sulphide: $\mathrm{Pb^{2+}(aq) + H_2S(g) \rightarrow PbS(s) + 2H^+(aq)}$ . Filter paper soaked in lead(II) ethanoate turns black in the gas, which is the usual test for hydrogen sulphide.

Worked example 1: comparing two acids

Estimate the pH of 0.10 mol dm⁻³ solutions of (a) hydrogen sulphide ( $\mathrm{p}K_{a1} = 7.0$ ) and (b) hydrogen selenide ( $\mathrm{p}K_{a1} = 3.89$ ), considering only the first dissociation. Explain the difference.

Solution. For a weak acid, $[\mathrm{H^+}] = \sqrt{K_a c}$ .
(a) $K_a = 10^{-7.0}$ : $[\mathrm{H^+}] = \sqrt{1.0 \times 10^{-7} \times 0.10} = 1.0 \times 10^{-4}\ \mathrm{mol\,dm^{-3}}$ , pH = 4.0.
(b) $K_a = 10^{-3.89} = 1.29 \times 10^{-4}$ : $[\mathrm{H^+}] = \sqrt{1.29 \times 10^{-4} \times 0.10} = 3.6 \times 10^{-3}\ \mathrm{mol\,dm^{-3}}$ , pH = 2.4.
The H–Se bond is longer and weaker than H–S, so $\mathrm{H_2Se}$ gives up its proton more easily: it is the stronger acid.

Worked example 2: making hydrogen sulphide

4.40 g of iron(II) sulphide reacts completely with excess dilute hydrochloric acid. Calculate the volume of hydrogen sulphide at room temperature and pressure (24.0 dm³ mol⁻¹). $A_{r}$ : Fe 55.85, S 32.07.

Solution. $\mathrm{FeS + 2HCl \rightarrow FeCl_2 + H_2S}$ : 1 mol FeS gives 1 mol $\mathrm{H_2S}$ . $M(\mathrm{FeS}) = 87.92\ \mathrm{g\,mol^{-1}} \qquad n = \frac{4.40}{87.92} = 0.0500\ \mathrm{mol} \qquad V = 0.0500 \times 24.0 = 1.20\ \mathrm{dm^3}$ That is enough gas to be deadly in a closed room, which is why the preparation is done only in a fume cupboard.

Warning

Common mistakes. (1) Expecting acid strength to follow electronegativity (" $\mathrm{H_2O}$ should be the strongest acid"): bond strength wins, so acidity increases down the group. (2) Explaining water's high boiling point by its "strong O–H bonds": boiling breaks hydrogen bonds between molecules, not the covalent bonds inside them. (3) Writing $\mathrm{H_2S + Cl_2}$ as giving $\mathrm{SCl_2}$ : chlorine oxidises $\mathrm{H_2S}$ to sulphur.

Safety. Hydrogen sulphide is highly toxic. Its smell is a warning at low levels, but at higher concentrations it quickly deadens the sense of smell, so a dangerous build-up may go unnoticed; NIOSH gives 100 ppm as immediately dangerous to life and health. It is prepared and used only in a fume cupboard. Anyone overcome by the gas should be moved to fresh air by rescuers who are themselves protected, and given medical help at once. Hydrogen selenide and hydrogen telluride are even more toxic.

Tip

Two opposite trends down Group 16: boiling point rises (London forces), but water is the odd one out (hydrogen bonds); acidity rises because the H–X bond weakens. Lower $\mathrm{p}K_a$ = stronger acid.

MSJ Chem: Group 16 hydrides (4 min). Watch for the explanation of the boiling-point trend from H₂S to H₂Te in terms of London forces, and why water is anomalous.

✦

Practice

State and explain the trend in boiling point from $\mathrm{H_2S}$ to $\mathrm{H_2Te}$ .
Explain why water boils at 100 °C, far above the boiling point of hydrogen sulphide.
Arrange $\mathrm{H_2O}$ , $\mathrm{H_2S}$ and $\mathrm{H_2Te}$ in order of increasing acid strength and explain the order.
Write equations for: (a) the laboratory preparation of hydrogen sulphide; (b) its complete combustion; (c) its reaction with chlorine.
Describe a test for hydrogen sulphide and write the equation.
Calculate the pH of 0.010 mol dm⁻³ hydrogen sulphide, using $\mathrm{p}K_{a1} = 7.0$ and ignoring the second dissociation.
In $\mathrm{2H_2S + SO_2 \rightarrow 3S + 2H_2O}$ , give the oxidation numbers of sulphur and identify the oxidising and reducing agents.
Predict, with a reason, whether $\mathrm{H_2Po}$ would be more or less thermally stable than $\mathrm{H_2Te}$ .

Show answersHide answers

It rises (−59.6, −41.3, −2.2 °C): more electrons give larger temporary dipoles and stronger London forces between molecules.
Water molecules form extensive hydrogen bonds (small, very electronegative O with two lone pairs); $\mathrm{H_2S}$ does not, because S is much less electronegative.
$\mathrm{H_2O}$ < $\mathrm{H_2S}$ < $\mathrm{H_2Te}$ : the H–X bond becomes longer and weaker down the group, so the proton is released more easily, and the larger anion holds it less tightly.
(a) $\mathrm{FeS + 2HCl \rightarrow FeCl_2 + H_2S}$ ; (b) $\mathrm{2H_2S + 3O_2 \rightarrow 2SO_2 + 2H_2O}$ ; (c) $\mathrm{H_2S + Cl_2 \rightarrow 2HCl + S}$ .
Filter paper moistened with lead(II) ethanoate turns black: $\mathrm{Pb^{2+} + H_2S \rightarrow PbS + 2H^+}$ .
$[\mathrm{H^+}] = \sqrt{10^{-7.0} \times 0.010} = 3.2 \times 10^{-5}$ ; pH = 4.5.
S is −2 in $\mathrm{H_2S}$ , +4 in $\mathrm{SO_2}$ and 0 in S. $\mathrm{SO_2}$ is the oxidising agent (reduced, +4 to 0); $\mathrm{H_2S}$ is the reducing agent (oxidised, −2 to 0).
Less stable: the H–Po bond would be even longer and weaker than H–Te.

Quick check

Which hydride is the strongest acid in water?

Key points

Boiling points rise from $\mathrm{H_2S}$ (−59.6 °C) to $\mathrm{H_2Te}$ (−2.2 °C) as London forces grow; water (100 °C) is anomalous because of hydrogen bonding.
Acid strength rises down the group ( $\mathrm{p}K_{a1}$ : $\mathrm{H_2S}$ 7.0, $\mathrm{H_2Se}$ 3.89, $\mathrm{H_2Te}$ 2.6) because the H–X bond becomes longer and weaker; thermal stability falls and reducing power rises.
$\mathrm{H_2S}$ is made from FeS and dilute HCl; it burns, reduces $\mathrm{Cl_2}$ , $\mathrm{SO_2}$ and $\mathrm{MnO_4^-}$ (giving sulphur), and precipitates metal sulphides; lead(II) ethanoate paper turns black.
$\mathrm{H_2S}$ is highly toxic and deadens the sense of smell: fume cupboard only.

Sources and credits

Image: Toestel-van-Kipp-L.jpg by Josh (nl.wikipedia), via Wikimedia Commons, licensed under CC BY-SA 3.0.
Data: boiling points and $\mathrm{p}K_a$ values from the Wikipedia articles Hydrogen sulfide, Hydrogen selenide and Hydrogen telluride (CRC Handbook values), CC BY-SA 4.0, adapted by Ganzaa.
Exposure limits: NIOSH Pocket Guide: Hydrogen sulfide, CDC.
Video: 4.4 Group 16 hydrides (SL), Mike Sugiyama Jones (MSJ Chem), embedded from YouTube.
The text, charts, table, worked examples and Practice items were written and drawn by Ganzaa.
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1Water and its poisonous cousin

2The data

3Volatility: London forces and one big exception

4Acidity: weaker bonds win

5Hydrogen sulphide in the laboratory

Water and its poisonous cousin

The data

Volatility: London forces and one big exception

Acidity: weaker bonds win

Hydrogen sulphide in the laboratory