Properties and Uses of Transition Metals

Lessons

Senior 6 · Chemistry

Properties and Uses of Transition Metals

5 lessons

Lesson 5 of 5

Identifying Transition Metal Ions, and What the Metals Are Used For

25 min

Which metal is in the bottle?

A laboratory technician finds an old bottle of pale green crystals with the label worn away. Is it an iron salt, a nickel salt or a chromium salt? Colour alone is not proof: several ions have similar shades, and the colour changes with the ligands present. Instead, chemists add two cheap reagents, sodium hydroxide solution and aqueous ammonia, and watch what forms. The colour of the precipitate, and whether it dissolves in excess reagent, identify the metal ion.

The tests

Put about 2 cm³ of the solution in a test tube. Add the reagent drop by drop, shaking gently, and note any precipitate. Then add the reagent in excess (about 5 cm³ more) and note whether the precipitate dissolves. Repeat with a fresh sample and the other reagent.

Table 5.1: Results with sodium hydroxide and ammonia solution

Ion	NaOH(aq), a few drops	NaOH(aq) in excess	NH₃(aq), a few drops	NH₃(aq) in excess
$\mathrm{Fe^{2+}}$	green precipitate; turns brown at the surface on standing	insoluble	green precipitate; turns brown on standing	insoluble
$\mathrm{Fe^{3+}}$	red-brown precipitate	insoluble	red-brown precipitate	insoluble
$\mathrm{Cu^{2+}}$	pale blue precipitate	insoluble	pale blue precipitate	dissolves: deep blue solution
$\mathrm{Mn^{2+}}$	off-white precipitate; darkens to brown in air	insoluble	off-white precipitate; darkens to brown	insoluble
$\mathrm{Cr^{3+}}$	grey-green precipitate	dissolves: dark green solution	grey-green precipitate	insoluble
$\mathrm{Zn^{2+}}$	white precipitate	dissolves: colourless solution	white precipitate	dissolves: colourless solution

Figure 5.1: The precipitates and solutions that identify five common ions. Sketch: real shades vary with concentration.

What is happening

Both reagents supply hydroxide ions. Sodium hydroxide is fully ionised; ammonia supplies a smaller concentration of $\mathrm{OH^-}$ through $\mathrm{NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)}$ . The hydroxide ions precipitate the insoluble metal hydroxides:

$\mathrm{Fe^{2+}(aq) + 2OH^-(aq) \rightarrow Fe(OH)_2(s)}$

$\mathrm{Fe^{3+}(aq) + 3OH^-(aq) \rightarrow Fe(OH)_3(s)}$

$\mathrm{Cu^{2+}(aq) + 2OH^-(aq) \rightarrow Cu(OH)_2(s)}$

At A-level you should also see this as a reaction of the aqua complex, in which hydroxide ions remove protons from two of the water ligands and leave a neutral, insoluble complex:

$\mathrm{[Fe(H_2O)_6]^{2+}(aq) + 2OH^-(aq) \rightarrow [Fe(OH)_2(H_2O)_4](s) + 2H_2O(l)}$

Green iron(II) hydroxide turns brown because oxygen from the air oxidises it to iron(III) hydroxide, starting at the surface. Recall from Lesson 1 that $\mathrm{Fe^{2+}}$ ( $3 d^{6}$ ) is readily oxidised to $\mathrm{Fe^{3+}}$ ( $3 d^{5}$ ):

$\mathrm{4Fe(OH)_2(s) + O_2(g) + 2H_2O(l) \rightarrow 4Fe(OH)_3(s)}$

In excess ammonia, copper(II) hydroxide dissolves because ammonia molecules become ligands, forming the deep blue complex from Lesson 3:

$\mathrm{Cu(OH)_2(s) + 4NH_3(aq) + 2H_2O(l) \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+}(aq) + 2OH^-(aq)}$

Zinc hydroxide is amphoteric: it dissolves in excess sodium hydroxide as the colourless zincate ion, $\mathrm{Zn(OH)_2(s) + 2OH^-(aq) \rightarrow [Zn(OH)_4]^{2-}(aq)}$ .

Definition

Amphoteric hydroxide: a hydroxide that reacts with both acids and bases. Example: zinc hydroxide dissolves in acids to give $\mathrm{Zn^{2+}}$ and in excess sodium hydroxide to give $\mathrm{[Zn(OH)_4]^{2-}}$ .

Confirmatory test: a second, more specific test used to confirm an ion suggested by a first test. Example: potassium thiocyanate turns $\mathrm{Fe^{3+}}$ solution blood-red.

Copper(II) sulfate solution with ammonia solution added. Most of the pale blue copper(II) hydroxide has dissolved as the deep blue ammine complex; a little undissolved precipitate remains at the bottom, so more ammonia is needed. Photo: Chemicalinterest, public domain, via Wikimedia Commons

Confirmatory tests for iron and copper

Table 5.2: Tests that tell the ions apart without hydroxide

Reagent	$\mathrm{Fe^{2+}}$	$\mathrm{Fe^{3+}}$	$\mathrm{Cu^{2+}}$
potassium thiocyanate, KSCN(aq)	no change (at most a faint pink)	deep blood-red solution, $\mathrm{[Fe(SCN)(H_2O)_5]^{2+}}$	no red colour
potassium hexacyanidoferrate(III), $\mathrm{K_3[Fe(CN)_6]}$	dark blue precipitate	brown solution, no precipitate	yellow-green precipitate
acidified potassium manganate(VII)	purple colour removed	no change	no change
potassium iodide, KI(aq)	no change	brown iodine forms	white precipitate (CuI) in brown iodine solution

The copper test is a redox reaction: $\mathrm{2Cu^{2+}(aq) + 4I^-(aq) \rightarrow 2CuI(s) + I_2(aq)}$ .

Worked example

Problem. Solution P is pale green. With sodium hydroxide solution it gives a green precipitate that does not dissolve in excess; after 20 minutes the top of the precipitate has turned orange-brown. Identify the cation in P, write ionic equations for both changes, and give one further test to confirm your answer.

Solution. A green hydroxide that is insoluble in excess and browns in air is iron(II) hydroxide, so P contains $\mathrm{Fe^{2+}}$ .

$\mathrm{Fe^{2+}(aq) + 2OH^-(aq) \rightarrow Fe(OH)_2(s)}$

$\mathrm{4Fe(OH)_2(s) + O_2(g) + 2H_2O(l) \rightarrow 4Fe(OH)_3(s)}$

Confirmation: add a few drops of acidified potassium manganate(VII); $\mathrm{Fe^{2+}}$ removes the purple colour. (Nickel(II), which is also green, gives a green hydroxide that does not turn brown and dissolves in excess ammonia to give a blue-violet solution.)

Worked example

Problem. Two unlabelled solutions are known to be iron(III) chloride and copper(II) sulfate. Using only aqueous ammonia, explain how to tell them apart, and write the equation for the change that identifies copper.

Solution. Add aqueous ammonia drop by drop, then in excess, to a sample of each. Iron(III) gives a red-brown precipitate of $\mathrm{Fe(OH)_3}$ that stays in excess. Copper(II) gives a pale blue precipitate that dissolves in excess to a deep blue solution:

$\mathrm{Cu(OH)_2(s) + 4NH_3(aq) + 2H_2O(l) \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+}(aq) + 2OH^-(aq)}$

Warning

Safety. Wear goggles: sodium hydroxide solution damages the eyes, even when dilute. Use ammonia solution in a well-ventilated room. Chromium(VI) compounds such as potassium dichromate are toxic and cause cancer; handle them with gloves. Do not pour solutions of copper, chromium, nickel or cobalt salts down the sink: collect them in the labelled waste bottle.

What the metals are used for

Table 5.3: Uses of the first-row transition metals, and the property behind each

Metal	Main uses	Property used
Titanium	aircraft parts, hip and knee replacements; $\mathrm{TiO_2}$ as the white pigment in paint and sunscreen	strong, low density, resists corrosion; $\mathrm{TiO_2}$ is white and scatters light
Vanadium	tough steels for tools and springs; $\mathrm{V_2O_5}$ catalyst in the contact process	hardens steel; variable oxidation state
Chromium	stainless steel; chrome plating	forms a thin, protective oxide layer
Manganese	steel for rails and hoes; $\mathrm{MnO_2}$ in dry cells	hardens steel; $\mathrm{MnO_2}$ is an oxidising agent
Iron	steel for buildings, bridges, roofing sheets and tools; Haber process catalyst	strong, cheap, easily alloyed
Cobalt	alloys for jet engines; lithium-ion battery cathodes; cobalt-60 in radiotherapy	keeps its strength when hot; variable oxidation state
Nickel	stainless steel; catalyst for hardening vegetable oils; rechargeable batteries	resists corrosion; adsorbs hydrogen on its surface
Copper	electrical wiring, water pipes; brass and bronze; copper fungicides such as Bordeaux mixture	excellent conductor, ductile, unreactive; $\mathrm{Cu^{2+}}$ is toxic to fungi
Zinc	galvanising steel roofing sheets; brass; dry-cell casings	more reactive than iron, so it corrodes first and protects it

Tip

Learn the three "dissolves in excess" results; everything else stays: zinc (both reagents, colourless), chromium(III) (NaOH only, dark green) and copper(II) (ammonia only, deep blue).

L Tull: Testing for positive metal ions, hydroxide precipitates (6 min). A real laboratory demonstration: watch the colours of the iron(II), iron(III) and copper(II) precipitates as sodium hydroxide is added.

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Practice

Give the colour of the precipitate formed when sodium hydroxide solution is added to solutions containing (a) $\mathrm{Fe^{2+}}$ , (b) $\mathrm{Fe^{3+}}$ , (c) $\mathrm{Cu^{2+}}$ .
Write ionic equations, with state symbols, for the reactions in question 1.
Explain why the precipitate from iron(II) turns brown at the surface when it is left to stand.
A solution gives a white precipitate with both reagents, and the precipitate dissolves in excess of both. Identify the cation and name the complex ion formed with excess ammonia.
Write the equation for the dissolving of copper(II) hydroxide in excess ammonia. What type of reaction is it?
Describe a test, other than with hydroxide ions, that gives a positive result with $\mathrm{Fe^{3+}}$ but not with $\mathrm{Fe^{2+}}$ .
A green solution gives a grey-green precipitate with sodium hydroxide, which dissolves in excess to a dark green solution. Identify the ion and suggest why a student who relied only on the colour of the original solution might have said iron(II).
A student adds excess sodium hydroxide to a solution containing both $\mathrm{Zn^{2+}}$ and $\mathrm{Cu^{2+}}$ , then filters. What is left on the filter paper, and what is in the filtrate? Explain.

Show answersHide answers

(a) green; (b) red-brown; (c) pale blue.
$\mathrm{Fe^{2+}(aq) + 2OH^-(aq) \rightarrow Fe(OH)_2(s)}$ ; $\mathrm{Fe^{3+}(aq) + 3OH^-(aq) \rightarrow Fe(OH)_3(s)}$ ; $\mathrm{Cu^{2+}(aq) + 2OH^-(aq) \rightarrow Cu(OH)_2(s)}$ .
Oxygen from the air oxidises iron(II) hydroxide to red-brown iron(III) hydroxide, starting where the precipitate touches the air: $\mathrm{4Fe(OH)_2 + O_2 + 2H_2O \rightarrow 4Fe(OH)_3}$ .
$\mathrm{Zn^{2+}}$ ; tetraamminezinc(II), $\mathrm{[Zn(NH_3)_4]^{2+}}$ .
$\mathrm{Cu(OH)_2(s) + 4NH_3(aq) + 2H_2O(l) \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+}(aq) + 2OH^-(aq)}$ . Ligand exchange (complex formation): ammonia molecules become ligands.
Add potassium thiocyanate solution: $\mathrm{Fe^{3+}}$ gives a deep blood-red colour; $\mathrm{Fe^{2+}}$ does not. (Or potassium iodide: $\mathrm{Fe^{3+}}$ oxidises iodide to brown iodine.)
$\mathrm{Cr^{3+}}$ . Chromium(III) and iron(II) solutions are both green, so colour alone cannot tell them apart; only chromium(III) hydroxide dissolves in excess sodium hydroxide.
Copper(II) hydroxide (pale blue) stays on the filter paper, because it is insoluble in excess sodium hydroxide. Zinc hydroxide is amphoteric and dissolves in excess as colourless $\mathrm{[Zn(OH)_4]^{2-}}$ , which passes into the filtrate.

1Which metal is in the bottle?

2The tests

3What is happening

4Confirmatory tests for iron and copper

5What the metals are used for

Which metal is in the bottle?

The tests

What is happening

Confirmatory tests for iron and copper

What the metals are used for