Reduction and Oxidation Reactions

Lessons

Senior 4 · Chemistry

Reduction and Oxidation Reactions

5 lessons

Lesson 5 of 5

The Reactivity Series of Metals

30 min

Why gold lies in the ground and sodium sits under oil

Gold is found as shining nuggets of the pure metal, unchanged after millions of years in wet soil. Sodium has to be stored under oil, and reacts with the moisture on your skin. Iron stands in between: it is extracted from its ores with carbon, and turns back into oxide (rust) if left in damp air. The reactivity series arranges metals in order of how readily they give up electrons, and so explains where they are found, how they react, how they are extracted, and how a more reactive metal can protect a less reactive one.

Definition

Reactivity series: a list of metals (with hydrogen and carbon included for reference) in order of decreasing tendency to lose electrons and form positive ions.

Displacement reaction: a redox reaction in which a more reactive metal takes the place of a less reactive metal in a compound. Example: $\mathrm{Fe(s) + Cu^{2+}(aq) \rightarrow Fe^{2+}(aq) + Cu(s)}$ .

Standard electrode potential, $E^\circ$ : the voltage of a half-cell, measured against the standard hydrogen electrode (defined as 0 V) at 25 °C, 1 atm and 1 mol dm⁻³. The more negative $E^\circ(\mathrm{M^{n+}/M})$ , the more easily the metal is oxidised.

The series and the evidence for it

Metal	With cold water	With dilute acid	Extracted by
potassium K	violent; hydrogen ignites (lilac flame)	explosive (never tried)	electrolysis of the molten compound
sodium Na	vigorous, melts into a ball	explosive (never tried)	electrolysis
calcium Ca	steady fizzing	vigorous	electrolysis
magnesium Mg	very slow; burns in steam	fast	electrolysis
aluminium Al	none (oxide layer); reacts with steam if powdered	slow at first, then fast	electrolysis of alumina
carbon C	reference point (not a metal)	—	metals below carbon: their oxides are reduced by heating with carbon
zinc Zn	none; reacts with steam	moderate	reduction with carbon
iron Fe	rusts slowly; reacts with steam	slow	reduction with carbon (blast furnace)
tin Sn, lead Pb	none	very slow	reduction with carbon
hydrogen H	reference point (not a metal)	metals below hydrogen give no hydrogen with dilute acids	—
copper Cu	none	none	reduction of sulfide ores; refined by electrolysis
silver Ag, gold Au	none	none	found native (uncombined) or extracted from ores

A metal high in the series loses electrons easily, so it is a strong reducing agent; its ions are hard to reduce, which is why the top metals must be extracted by electrolysis. The order matches the standard electrode potentials (Figure 16.5): the more negative $E^\circ$ , the stronger the metal as a reducing agent.

Figure 16.5: Standard electrode potentials of metals, from Wikipedia's data page. Note that calcium (−2.87 V) is more negative than sodium (−2.71 V), yet sodium reacts more violently with water: E° measures how far a reaction tends to go, not how fast it goes.

Displacement reactions

A metal displaces any metal below it from a solution of its salt. Iron in copper(II) sulfate becomes coated with copper and the blue colour fades: $\mathrm{Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)} \qquad \text{ionic: } \mathrm{Fe + Cu^{2+} \rightarrow Fe^{2+} + Cu}$ Copper does not react with iron(II) sulfate, because copper is below iron. A metal also displaces a less reactive metal from its oxide when heated. In the thermite reaction, used to weld railway lines, aluminium reduces iron(III) oxide so exothermically that the iron melts: $\mathrm{2Al(s) + Fe_2O_3(s) \rightarrow Al_2O_3(s) + 2Fe(l)}$

Predicting with electrode potentials

For two metals joined in a cell, the more negative metal is the anode (it is oxidised) and $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ A positive $E^\circ_{\text{cell}}$ means the reaction can happen. The further apart two metals are in the series, the larger the voltage of a cell made from them. Magnesium and copper: $E^\circ_{\text{cell}} = (+0.337) - (-2.372) = +2.71\ \mathrm{V}$ .

Using the series: sacrificial protection

A more reactive metal in contact with iron is oxidised in its place. Steel is coated with zinc (galvanised) for roofing sheets and buckets, and blocks of magnesium or zinc are bolted to ship hulls and buried pipelines; the block corrodes away while the steel is protected.

A gold nugget: crystals of the uncombined metal, as found in the ground. Gold is at the bottom of the reactivity series (E° = +1.52 V for Au³⁺/Au), so it does not react with water, air or dilute acids. Photo: United States Government, public domain, via Wikimedia Commons

Worked example 1: predicting reactions

Use Figure 16.5 to predict whether a reaction occurs, and write an ionic equation where it does: (a) zinc added to lead(II) nitrate solution; (b) silver added to dilute hydrochloric acid; (c) copper added to silver nitrate solution.

Solution. (a) Zn (−0.76 V) is above Pb (−0.13 V): $\mathrm{Zn + Pb^{2+} \rightarrow Zn^{2+} + Pb}$ ; $E^\circ_{\text{cell}} = -0.13 - (-0.76) = +0.63\ \mathrm{V}$ . (b) Ag (+0.80 V) is below hydrogen (0 V): $E^\circ_{\text{cell}} = 0 - 0.80 = -0.80\ \mathrm{V}$ , negative, so no reaction. (c) Cu (+0.34 V) is above Ag (+0.80 V): $\mathrm{Cu + 2Ag^+ \rightarrow Cu^{2+} + 2Ag}$ ; $E^\circ_{\text{cell}} = 0.80 - 0.34 = +0.46\ \mathrm{V}$ .

Worked example 2: copper from iron nails

Iron nails are left in excess copper(II) sulfate solution until 2.80 g of iron has dissolved as iron(II) ions. Calculate the mass of copper deposited. (Fe = 55.85, Cu = 63.55)

Solution. $\mathrm{Fe + Cu^{2+} \rightarrow Fe^{2+} + Cu}$ : 1 mol Fe gives 1 mol Cu. $n(\mathrm{Fe}) = \frac{2.80}{55.85} = 5.01 \times 10^{-2}\ \mathrm{mol} \qquad m(\mathrm{Cu}) = 5.01 \times 10^{-2} \times 63.55 = 3.19\ \mathrm{g}$ The solid gains mass (3.19 g of copper replaces 2.80 g of iron), a useful check in the laboratory. Recovering copper from waste solutions with scrap iron in this way is called cementation.

Warning

Common mistakes. (1) Saying a metal "displaces" one above it: only a more reactive metal displaces a less reactive one. (2) Thinking aluminium is unreactive because pans do not corrode: its oxide layer hides a reactive metal. (3) Putting carbon and hydrogen into the series as if they were metals: they are reference points. (4) Using $E^\circ$ to predict how fast a reaction is: it tells you whether a reaction can happen, not its rate.

Safety. Potassium, sodium and calcium are demonstrated by the teacher only, in tiny pieces behind a safety screen. Do not add potassium or sodium to acids. The thermite reaction throws out molten iron and is never done in a school laboratory. Hydrogen from metals and acids forms explosive mixtures with air: keep flames away except for a small test-tube "pop" test.

Tip

"Please Send Cats, Monkeys And Cute Zebras Into Safe Lovely Hot Countries, Signing Gold": K, Na, Ca, Mg, Al, (C), Zn, Fe, Sn, Pb, (H), Cu, Ag, Au. Higher displaces lower; above C needs electrolysis; below H ignores dilute acid.

Royal Society of Chemistry: Reactivity Series of Metals (16 min). Watch the alkali metals with water, metals with acid, and displacement reactions, and pause to predict each result before it is shown.

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Practice

Arrange these metals in order of decreasing reactivity: copper, magnesium, iron, sodium, zinc.
Predict, with a reason, whether a reaction occurs, and write an ionic equation if it does: (a) magnesium in zinc sulfate solution; (b) copper in iron(II) sulfate solution; (c) iron in dilute sulfuric acid.
Explain why aluminium is extracted by electrolysis but iron by reduction with carbon.
Calculate $E^\circ_{\text{cell}}$ for a cell made from zinc and silver half-cells, and state which metal is the anode.
Explain why a zinc coating protects steel even when the coating is scratched.
0.486 g of magnesium is added to excess silver nitrate solution. Calculate the mass of silver formed (Mg = 24.31, Ag = 107.87).
Use Figure 16.5 to explain why sodium is placed above calcium in the reactivity series, even though $E^\circ$ for calcium is more negative.
A student places strips of metals X, Y and Z in solutions of each other's nitrates. X displaces both Y and Z; Z displaces Y. Put X, Y and Z in order of reactivity.

Show answersHide answers

Sodium, magnesium, zinc, iron, copper.
(a) Yes: Mg is above Zn; $\mathrm{Mg + Zn^{2+} \rightarrow Mg^{2+} + Zn}$ . (b) No: Cu is below Fe. (c) Yes: Fe is above H; $\mathrm{Fe + 2H^+ \rightarrow Fe^{2+} + H_2}$ .
Aluminium is above carbon: carbon cannot remove oxygen from $\mathrm{Al_2O_3}$ , so electrolysis must be used. Iron is below carbon, so carbon (as coke and CO) reduces iron oxide.
$E^\circ_{\text{cell}} = 0.80 - (-0.76) = +1.56\ \mathrm{V}$ ; zinc (more negative) is the anode.
Zinc is more reactive than iron, so where both are exposed the zinc is oxidised instead of the iron: it is sacrificed, giving electrons to the iron.
$\mathrm{Mg + 2Ag^+ \rightarrow Mg^{2+} + 2Ag}$ ; $n(\mathrm{Mg}) = \frac{0.486}{24.31} = 0.0200\ \mathrm{mol}$ ; $n(\mathrm{Ag}) = 0.0400\ \mathrm{mol}$ ; mass $= 0.0400 \times 107.87 = 4.31\ \mathrm{g}$ .
The reactivity series is based on how vigorously metals react, which also depends on rate. Sodium reacts faster with water than calcium (it needs to lose only one electron, and calcium hydroxide is only slightly soluble), although $E^\circ$ shows calcium has the slightly greater tendency to be oxidised in solution.
X > Z > Y.

Quick check

Which metal lies above hydrogen but below carbon in the reactivity series, so it reacts with dilute acid and can be extracted with carbon?

Key points

Reactivity series: K, Na, Ca, Mg, Al, (C), Zn, Fe, Sn, Pb, (H), Cu, Ag, Au. Higher metals lose electrons more easily and are stronger reducing agents.
A metal displaces a lower metal from its salts or oxides; metals above hydrogen give hydrogen with dilute acids.
The order follows $E^\circ(\mathrm{M^{n+}/M})$ ; $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ , and a positive value means the reaction can occur.
Position decides extraction (electrolysis above carbon, carbon reduction below it, native metals at the bottom) and allows sacrificial protection of iron by zinc or magnesium.

1Why gold lies in the ground and sodium sits under oil

2The series and the evidence for it

3Displacement reactions

4Predicting with electrode potentials

5Using the series: sacrificial protection

Why gold lies in the ground and sodium sits under oil

The series and the evidence for it

Displacement reactions

Predicting with electrode potentials

Using the series: sacrificial protection