Trends in Chemical Properties of Group 13 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 13 Elements and Their Compounds

4 lessons

Lesson 3 of 4

Amphoteric Oxides and Hydroxides, and the Test for Aluminium Ions

25 min

Separating aluminium from rust-coloured rock

Bauxite is red-brown because, besides hydrated aluminium oxide, it contains iron(III) oxide, silica and titanium dioxide. To make pure aluminium, the alumina must first be separated from all of these, and the method that does it, used at every alumina refinery in the world, depends on one property: aluminium oxide dissolves in hot sodium hydroxide solution, while iron(III) oxide does not. An oxide that reacts with acids and with alkalis is called amphoteric. This lesson explains why aluminium and gallium form amphoteric oxides and hydroxides, and how the same property gives a simple test for aluminium ions.

Definition

Amphoteric: describes an oxide or hydroxide that reacts both with acids, acting as a base, and with alkalis, acting as an acid. Example: $\mathrm{Al(OH)_3}$ dissolves in hydrochloric acid to give $\mathrm{Al^{3+}(aq)}$ , and in sodium hydroxide solution to give $\mathrm{[Al(OH)_4]^-(aq)}$ .

Aluminate ion: $\mathrm{[Al(OH)_4]^-}$ , tetrahydroxoaluminate(III), the soluble form of aluminium in alkaline solution. The gallium analogue is $\mathrm{[Ga(OH)_4]^-}$ , gallate.

Hydrolysis of a cation: the reaction of a hydrated metal ion with water in which a water ligand loses a proton, making the solution acidic. Example: $\mathrm{[Al(H_2O)_6]^{3+} + H_2O \rightleftharpoons [Al(H_2O)_5(OH)]^{2+} + H_3O^+}$ .

Aluminium oxide and hydroxide: two faces

As a base, with acids:

\mathrm{Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)}

\mathrm{Al(OH)_3(s) + 3H^+(aq) \rightarrow Al^{3+}(aq) + 3H_2O(l)}

As an acid, with alkalis:

\mathrm{Al_2O_3(s) + 2NaOH(aq) + 3H_2O(l) \rightarrow 2Na[Al(OH)_4](aq)}

\mathrm{Al(OH)_3(s) + OH^-(aq) \rightarrow [Al(OH)_4]^-(aq)}

Gallium(III) oxide and hydroxide behave in the same way, for example $\mathrm{Ga(OH)_3 + OH^- \rightarrow [Ga(OH)_4]^-}$ . Strongly heated (calcined) aluminium oxide, such as corundum, reacts only slowly with both, because its lattice is very strong; freshly precipitated aluminium hydroxide reacts quickly.

Why aluminium is amphoteric

In water, $\mathrm{Al^{3+}}$ is surrounded by six water molecules, $\mathrm{[Al(H_2O)_6]^{3+}}$ . The small, highly charged ion pulls electron density from the O–H bonds of these water ligands, so they lose protons easily. The hexaaqua ion is a weak acid, with $K_a \approx 10^{-5}\ \mathrm{mol\,dm^{-3}}$ , about as strong as ethanoic acid, which is why solutions of aluminium salts are acidic.

Adding hydroxide ions removes protons one at a time (Figure 8.4). When three have gone, the neutral $\mathrm{[Al(H_2O)_3(OH)_3]}$ , aluminium hydroxide, comes out of solution as a white, jelly-like precipitate. With excess hydroxide a fourth proton is removed and the charged, soluble aluminate ion forms. Adding acid reverses every step. Amphoteric behaviour is simply this ladder being climbed in either direction.

Figure 8.4: Each step down removes one $\mathrm{H^+}$ from a water ligand. Only the neutral species, $\mathrm{Al(OH)_3}$ , is insoluble.

The trend down Group 13

Table 8.4: Acid–base character of the oxides and hydroxides

Element	Oxide	Hydroxide	Character
Boron	$\mathrm{B_2O_3}$	$\mathrm{B(OH)_3}$ (boric acid)	acidic (weak acid)
Aluminium	$\mathrm{Al_2O_3}$	$\mathrm{Al(OH)_3}$	amphoteric
Gallium	$\mathrm{Ga_2O_3}$	$\mathrm{Ga(OH)_3}$	amphoteric
Indium	$\mathrm{In_2O_3}$	$\mathrm{In(OH)_3}$	basic; does not dissolve in alkali
Thallium	$\mathrm{Tl_2O}$ (+1)	TlOH (+1)	strongly basic; TlOH is soluble

The oxides become more basic down the group, as the elements become more metallic and their cations larger and less polarising. The pattern is the same one you met in Period 3: a non-metal oxide is acidic, a metal oxide basic, and an element on the borderline gives an amphoteric oxide.

Using amphoterism: the Bayer process

Digestion. Crushed bauxite is heated with concentrated sodium hydroxide solution at 150–200 °C under pressure. The alumina dissolves: $\mathrm{Al_2O_3 + 2OH^- + 3H_2O \rightarrow 2[Al(OH)_4]^-}$ . Basic iron(III) oxide and titanium dioxide do not dissolve.
Separation. The undissolved solid, red mud, is filtered off.
Precipitation. The solution is cooled and seeded with crystals of aluminium hydroxide, which grows as pure $\mathrm{Al(OH)_3}$ : $\mathrm{[Al(OH)_4]^- \rightarrow Al(OH)_3 + OH^-}$ .
Calcination. The hydroxide is heated to about 1200 °C: $\mathrm{2Al(OH)_3 \rightarrow Al_2O_3 + 3H_2O}$ . The pure alumina goes to the Hall–Héroult cell (Lesson 1).

Because gallium is chemically so like aluminium, it dissolves along with it; it builds up in the recycled liquor, which is the world's main source of gallium.

Aerial view of the bauxite residue (red mud) storage area at the Aughinish alumina refinery in Ireland, 2010. The red colour is iron(III) oxide, the basic oxide left behind when the amphoteric alumina dissolves in sodium hydroxide. Photo: Graham Horn, CC BY-SA 2.0, via Wikimedia Commons (Geograph)

Testing for aluminium ions

Solutions of aluminium salts are colourless, and aluminium gives no colour in a flame test, so it is identified with hydroxide precipitates.

Sodium hydroxide solution, drop by drop, then in excess. A white gelatinous precipitate forms, then dissolves in excess to give a colourless solution (Figure 8.5): $\mathrm{Al^{3+}(aq) + 3OH^-(aq) \rightarrow Al(OH)_3(s)} \qquad \mathrm{Al(OH)_3(s) + OH^-(aq) \rightarrow [Al(OH)_4]^-(aq)}$
Aqueous ammonia, drop by drop, then in excess. The same white precipitate forms but does not dissolve in excess, because ammonia is a weak base and never makes the solution alkaline enough to form the aluminate ion.

Figure 8.5: The sodium hydroxide test for $\mathrm{Al^{3+}}$ .

Zinc and lead(II) ions also give white precipitates that dissolve in excess sodium hydroxide, so one test alone is not enough. Use the flowchart in Figure 8.6: zinc hydroxide dissolves in excess ammonia (forming $\mathrm{[Zn(NH_3)_4]^{2+}}$ ), and lead(II) gives a yellow precipitate of lead(II) iodide with potassium iodide; aluminium does neither.

Figure 8.6: Identifying a white hydroxide precipitate.

Worked example 1: how much alkali?

Sodium hydroxide solution, 0.500 mol dm⁻³, is added from a burette to 25.0 cm³ of 0.100 mol dm⁻³ aluminium chloride solution. Calculate (a) the volume added when the mass of precipitate is greatest, (b) that mass, and (c) the total volume needed to dissolve the precipitate completely. $M(\mathrm{Al(OH)_3}) = 78.0\ \mathrm{g\,mol^{-1}}$ .

Solution. $n(\mathrm{Al^{3+}}) = 0.0250\ \mathrm{dm^3} \times 0.100\ \mathrm{mol\,dm^{-3}} = 2.50 \times 10^{-3}\ \mathrm{mol}$ .

(a) $\mathrm{Al^{3+} + 3OH^- \rightarrow Al(OH)_3}$ needs $3 \times 2.50 \times 10^{-3} = 7.50 \times 10^{-3}$ mol $\mathrm{OH^-}$ : $V = \frac{7.50 \times 10^{-3}}{0.500} = 0.0150\ \mathrm{dm^3} = 15.0\ \mathrm{cm^3}$ .
(b) Mass $= 2.50 \times 10^{-3} \times 78.0 = 0.195\ \mathrm{g}$ .
(c) $\mathrm{Al(OH)_3 + OH^- \rightarrow [Al(OH)_4]^-}$ needs a further $2.50 \times 10^{-3}$ mol, which is 5.00 cm³. Total $= 15.0 + 5.00 = 20.0\ \mathrm{cm^3}$ .

Worked example 2: why aluminium salts are acidic

Estimate the pH of 0.100 mol dm⁻³ aluminium chloride solution, taking $K_a = 1.0 \times 10^{-5}\ \mathrm{mol\,dm^{-3}}$ for $\mathrm{[Al(H_2O)_6]^{3+}}$ and treating only the first step.

Solution. As for any weak acid, assuming little dissociation: $[\mathrm{H_3O^+}] = \sqrt{K_a \times c} = \sqrt{1.0 \times 10^{-5} \times 0.100} = 1.0 \times 10^{-3}\ \mathrm{mol\,dm^{-3}}$ $\mathrm{pH} = -\log_{10}(1.0 \times 10^{-3}) = 3.0$ So the solution is distinctly acidic even though no acid was added: blue litmus turns red, and adding sodium carbonate gives bubbles of carbon dioxide.

Warning

Common mistakes. (1) Saying the precipitate dissolves in excess ammonia: aluminium hydroxide does not; zinc hydroxide does. (2) Concluding "aluminium" from the sodium hydroxide test alone: zinc and lead(II) behave the same way. (3) Writing the aluminate ion as $\mathrm{AlO_2^-}$ in a solution equation: in water it is $\mathrm{[Al(OH)_4]^-}$ . (4) Adding one large portion of alkali at once: the precipitate may form and redissolve before you see it, so add it drop by drop.

Safety. Sodium hydroxide solution is corrosive, even at 0.5 mol dm⁻³, and can blind: wear goggles. Aqueous ammonia gives off an irritating gas, so use it in a well-ventilated room. Lead compounds are toxic: use dilute solutions, avoid skin contact and wash your hands after the practical. If a solution splashes into your eyes, rinse them at once with plenty of running water for several minutes and get medical help.

Tip

Remember the aluminium test as "NaOH: comes and goes; NH₃: comes and stays." Then rule out the look-alikes: zinc goes in ammonia, lead turns yellow with iodide.

Chemistry and Biochemistry Demo Lab at Oregon State University: Amphoteric Hydroxides (3 min). Watch aluminium hydroxide and iron(III) hydroxide being treated with acid and with excess base: both dissolve in acid, but only aluminium hydroxide dissolves in excess base.

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Practice

Define amphoteric and write two ionic equations that show aluminium hydroxide is amphoteric.
Write an equation for the reaction of gallium(III) oxide with sodium hydroxide solution.
Arrange $\mathrm{B_2O_3}$ , $\mathrm{Al_2O_3}$ and $\mathrm{Tl_2O}$ in order of increasing basic character, and explain the trend.
Explain why a solution of aluminium sulphate turns blue litmus red.
In the Bayer process, why is iron(III) oxide left behind when bauxite is digested with sodium hydroxide solution?
A colourless solution gives a white precipitate with a few drops of aqueous ammonia; the precipitate dissolves in excess ammonia. Which ion from Figure 8.6 is present?
20.0 cm³ of a solution of aluminium nitrate needed exactly 24.0 cm³ of 0.250 mol dm⁻³ sodium hydroxide to give the greatest mass of precipitate. Calculate the concentration of the aluminium nitrate, and the further volume of the alkali that would dissolve the precipitate.

Show answersHide answers

Amphoteric: reacts with both acids and alkalis. $\mathrm{Al(OH)_3 + 3H^+ \rightarrow Al^{3+} + 3H_2O}$ ; $\mathrm{Al(OH)_3 + OH^- \rightarrow [Al(OH)_4]^-}$ .
$\mathrm{Ga_2O_3 + 2NaOH + 3H_2O \rightarrow 2Na[Ga(OH)_4]}$ .
$\mathrm{B_2O_3}$ (acidic) < $\mathrm{Al_2O_3}$ (amphoteric) < $\mathrm{Tl_2O}$ (strongly basic). Down the group the elements become more metallic and the cations larger and less polarising, so the oxides become more ionic and more basic.
The hydrated ion $\mathrm{[Al(H_2O)_6]^{3+}}$ is a weak acid: the small, highly charged $\mathrm{Al^{3+}}$ weakens the O–H bonds of its water ligands, which give protons to water, forming $\mathrm{H_3O^+}$ .
Iron(III) oxide is basic, not amphoteric, so it does not react with the alkali; only the amphoteric alumina dissolves.
$\mathrm{Zn^{2+}}$ .
$n(\mathrm{OH^-}) = 0.0240 \times 0.250 = 6.00 \times 10^{-3}$ mol; $n(\mathrm{Al^{3+}}) = \frac{6.00 \times 10^{-3}}{3} = 2.00 \times 10^{-3}$ mol; concentration $= \frac{2.00 \times 10^{-3}}{0.0200} = 0.100\ \mathrm{mol\,dm^{-3}}$ . A further $2.00 \times 10^{-3}$ mol of $\mathrm{OH^-}$ is needed: $\frac{2.00 \times 10^{-3}}{0.250} = 8.00\ \mathrm{cm^3}$ .

Quick check

A colourless solution gives a white precipitate with sodium hydroxide that dissolves in excess. With aqueous ammonia the white precipitate stays in excess, and potassium iodide gives no precipitate. Which ion is present?

Key points

$\mathrm{Al_2O_3}$ , $\mathrm{Al(OH)_3}$ , $\mathrm{Ga_2O_3}$ and $\mathrm{Ga(OH)_3}$ are amphoteric: they give $\mathrm{M^{3+}}$ with acids and $\mathrm{[M(OH)_4]^-}$ with alkalis. Down Group 13 the oxides change from acidic (B) to amphoteric (Al, Ga) to basic (In, Tl).
$\mathrm{[Al(H_2O)_6]^{3+}}$ is a weak acid ( $K_a \approx 10^{-5}$ ), so aluminium salt solutions are acidic; removing protons step by step gives $\mathrm{Al(OH)_3}$ , then $\mathrm{[Al(OH)_4]^-}$ .
The Bayer process uses amphoterism to separate alumina from basic iron(III) oxide (red mud).
Test for $\mathrm{Al^{3+}}$ : white precipitate with NaOH(aq), soluble in excess; white precipitate with NH₃(aq), insoluble in excess; no yellow precipitate with KI(aq).

Sources and credits

Image: Aughinish West - geograph.org.uk - 1873607.jpg by Graham Horn, via Wikimedia Commons, licensed under CC BY-SA 2.0.
Image: Potassium alum octahedral crystal.jpg by Ude, via Wikimedia Commons, licensed under CC BY-SA 3.0.
Facts on the hexaaqua ion and its $K_{a}$ : Wikipedia, Aluminium; Bayer process conditions and gallium recovery: Bayer process; indium's insolubility in alkali: Indium; all CC BY-SA 4.0, adapted by Ganzaa.
First aid for chemical splashes: NIOSH Pocket Guide, first aid, CDC.
Video: Amphoteric Hydroxides, Chemistry and Biochemistry Demo Lab at OSU, embedded from YouTube.
The text, diagrams, flowchart, 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.

1Separating aluminium from rust-coloured rock

2Aluminium oxide and hydroxide: two faces

3Why aluminium is amphoteric

4The trend down Group 13

5Using amphoterism: the Bayer process

6Testing for aluminium ions