Chemistry · Solutions and electrolytes · Grade 9

Hydrolysis of salts and the medium

Why is a soda solution alkaline, ammonium chloride acidic and table salt neutral?

The situation

My mother dissolved washing soda in hot water and washed the dishes — the skin of her hands became slippery, as if from soap. A craftsman cleans metal with a solution of sal ammoniac (ammonium chloride) before soldering, and its solution is acidic. A solution of table salt, on the other hand, is neither alkaline nor acidic. All three substances are salts: they are neither an acid nor an alkali. So why do their solutions give different media, and why does hot soda wash better?

Topic

Hydrolysis of salts and the reaction of a solution. Washing soda makes hands feel “soapy”, ammonium chloride is acidic, table salt is neutral — the pH of a salt solution

Key ideas

Hydrolysis by cation and anion
pH of a salt solution
Degree of hydrolysis

Formulas

CO_{3}^{2 -} + H_{2} O \overset{}{⇌} HCO_{3}^{-} + OH^{-}

Anion hydrolysis: the anion of a weak acid takes H⁺ from water — OH⁻ is left in excess, and the medium is alkaline.

NH_{4}^{+} + H_{2} O \overset{}{⇌} NH_{3} \cdot H_{2} O + H^{+}

Cation hydrolysis: the cation of a weak base gives H⁺ to water — the medium is acidic.

K_{g} = \frac{K_{w}}{K_{a}}, h \approx \sqrt{\frac{K_{g}}{c}}

The hydrolysis constant and degree of hydrolysis (for an anion; for a cation $K_{g} = K_{w} / K_{b}$ ). On dilution and heating h increases.

pH = 7 + \frac{1}{2} p K_{a} + \frac{1}{2} \lg c

The pH of a solution of a salt of a weak acid and a strong base (25 °C, approximate formula).

2 Al^{3 +} + 3 CO_{3}^{2 -} + 3 H_{2} O \overset{}{\to} 2 Al (OH)_{3} ↓ + 3 CO_{2} ↑

Complete (joint) hydrolysis: when washing soda is added to an aluminium salt, a precipitate of Al(OH)₃ and CO₂ gas form.

Theory

What hydrolysis is and when it occurs

Water itself contains very few H⁺ and OH⁻ ions: $H_{2} O \overset{}{⇌} H^{+} + OH^{-}$ , and at 25 °C $[H^{+}] [O H^{-}] = K_{w} = 10^{- 14}$ . When a salt dissolves, its ions sometimes form a weak electrolyte with the ions of water — the equilibrium is disturbed and H⁺ or OH⁻ is left in excess in the solution. This exchange reaction of the ions of a salt with water is hydrolysis. Depending on which base and which acid formed the salt, there are 4 cases: strong base + strong acid (NaCl, KNO₃, K₂SO₄) — no hydrolysis, pH ≈ 7; strong base + weak acid (Na₂CO₃, CH₃COONa, K₂SiO₃) — anion hydrolysis, the medium is alkaline; weak base + strong acid (NH₄Cl, AlCl₃, ZnCl₂, CuSO₄) — cation hydrolysis, the medium is acidic; weak base + weak acid (CH₃COONH₄, (NH₄)₂CO₃) — hydrolysis of both cation and anion, the medium depends on $K_{a}$ and $K_{b}$ (almost neutral in ammonium ethanoate). Hydrolysis is the reverse of neutralisation and usually proceeds only to a small extent — most of the salt ions in the solution remain unchanged.

Anion hydrolysis

In a soda solution the carbonate ion takes a proton from a water molecule: $CO_{3}^{2 -} + H_{2} O \overset{}{⇌} HCO_{3}^{-} + OH^{-}$ , in molecular form: $Na_{2} CO_{3} + H_{2} O \overset{}{⇌} NaHCO_{3} + NaOH$ . Highly charged anions hydrolyse in stages, but in practice only the first stage matters: the second ( $HCO_{3}^{-} + H_{2} O \overset{}{⇌} H_{2} CO_{3} + OH^{-}$ ) is millions of times weaker than the first. In a 0.1 mol/L solution of Na₂CO₃ only ≈ 4.5 % of the carbonate ions are hydrolysed, but the OH⁻ formed ( $\approx 4.5 \cdot 10^{- 3}$ mol/L) raises the pH to ≈ 11.7 — which is why soda removes grease well and affects the skin. Sodium ethanoate (pH ≈ 8.9) and sodium phosphate (pH ≈ 12.6) behave in the same way. In baking soda, NaHCO₃, the HCO₃⁻ ion both accepts and donates a proton; accepting a proton (hydrolysis) is stronger than donating, so the medium is weakly alkaline — pH ≈ 8.3.

Cation hydrolysis. Aluminium salts (lab work 11)

The ammonium ion gives its proton to water: $NH_{4}^{+} + H_{2} O \overset{}{⇌} NH_{3} \cdot H_{2} O + H^{+}$ — in 0.1 mol/L NH₄Cl the pH ≈ 5.1. Highly charged cations exist in water as aqua complexes, and their water ligands readily give up a proton: $[Al (H_{2} O)_{6}]^{3 +} + H_{2} O \overset{}{⇌} [Al (OH) (H_{2} O)_{5}]^{2 +} + H_{3} O^{+}$ — so in a 0.1 mol/L solution of AlCl₃ the pH ≈ 3.0: litmus turns red, methyl orange turns red-pink and phenolphthalein stays colourless. Lab work: add drops of indicators to solutions of AlCl₃, Al₂(SO₄)₃ and alum and compare them with the control (distilled water). If soda is added to an aluminium salt, the two hydrolyses reinforce each other — H⁺ and OH⁻ combine to form water and the equilibria shift all the way: $2 Al^{3 +} + 3 CO_{3}^{2 -} + 3 H_{2} O \overset{}{\to} 2 Al (OH)_{3} ↓ + 3 CO_{2} ↑$ — a white gelatinous precipitate of Al(OH)₃ and bubbles of CO₂. This is why aluminium carbonate does not exist in aqueous solution. Clarifying water with alum is also based on the hydrolysis of Al³⁺.

The degree of hydrolysis and the factors that affect it

The degree of hydrolysis $h$ is the ratio of the number of hydrolysed ions to their total number. Dilution increases hydrolysis: if soda is diluted 10 times, h rises from ≈ 4.5 % to ≈ 13.6 %, but the total concentration of OH⁻ falls — the pH drops from 11.7 to 11.1. Heating also increases hydrolysis: hydrolysis is the reverse of neutralisation (ΔH ≈ −56 kJ/mol), that is, it is endothermic. A hot soda solution contains more OH⁻ — it washes better. If a solution of iron(III) chloride is heated, the yellow colour changes to reddish-brown: colloidal particles of $Fe (OH)_{3}$ form. To suppress hydrolysis, one of the products must be added: a little acid is added to solutions of FeCl₃ and AlCl₃, and an alkali to a soda solution (Le Chatelier's principle).

University: the hydrolysis constant

For the anion hydrolysis equilibrium $A^{-} + H_{2} O \overset{}{⇌} HA + OH^{-}$ , $K_{g} = \frac{[HA] [O H^{-}]}{[A^{-}]} = \frac{K_{w}}{K_{a}}$ ; for a cation $K_{g} = K_{w} / K_{b}$ , and for both ions $K_{g} = K_{w} / (K_{a} K_{b})$ . For small $h$ , $h \approx \sqrt{K_{g} / c}$ and $[O H^{-}] \approx \sqrt{K_{g} c}$ , from which $pH = 7 + \frac{1}{2} p K_{a} + \frac{1}{2} \lg c$ (for a cation $pH = 7 - \frac{1}{2} p K_{b} - \frac{1}{2} \lg c$ ; for ammonium ethanoate $pH = 7 + \frac{1}{2} (p K_{a} - p K_{b})$ — independent of concentration). On the graph, $\lg h$ against $\lg c$ is a straight line with slope −½; at high $h$ the exact solution departs from it. Temperature: from 25 to 50 °C $K_{w}$ increases ≈ 5.5 times while $K_{a}$ barely changes — $K_{g}$ and $h$ increase sharply. Interestingly, even though a hot soda solution contains more OH⁻, a pH meter shows a lower pH, because on heating the neutral point itself lies below pH 7. Complete hydrolysis: the salt of a weak base and a weak, volatile or sparingly soluble acid decomposes completely in water: $Al_{2} S_{3} + 6 H_{2} O \overset{}{\to} 2 Al (OH)_{3} ↓ + 3 H_{2} S ↑$ .

Safety and the model

Safety: solutions of soda and sodium phosphate are alkaline — they dry the skin; if they get into the eyes, rinse with plenty of water; work in safety goggles and gloves. Heated ammonium chloride releases ammonia and HCl — work in a fume cupboard. Put out a spirit burner with its cap, and never point the mouth of a test tube being heated at yourself or others. Model: the solutions are ideal (activities equal concentrations); the pH is calculated exactly from the charge-balance equation; $K_{a}$ values are from the CRC handbook, and for ${[Al (H_{2} O)_{6}]}^{3 +}$ $\lg K = - 4.97$ (Baes and Mesmer), first stage only; the formation of $HSO_{4}^{-}$ is taken into account in sulfates, while ion pairs and polynuclear species are not. Temperature dependence (0–60 °C): for $K_{w}$ and $K_{a}$ the equations of Harned and co-workers, for NH₄⁺ those of Bates and Pinching. The colour change of FeCl₃ on heating is shown qualitatively. The molecule scene is a diagram of one act of hydrolysis; in a real solution this process is very fast and is repeated many times.

Did you know?

Alum — KAl(SO₄)₂·12H₂O — has been used since antiquity to clarify water: flakes of Al(OH)₃, formed by the hydrolysis of Al³⁺ ions, stick to the particles of turbidity and settle.
The pH of human blood is kept between 7.35 and 7.45; the main role in this is played by the hydrogencarbonate buffer system (HCO₃⁻ / H₂CO₃).

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

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