Reduction and Oxidation Reactions

Lessons

Senior 4 · Chemistry

Reduction and Oxidation Reactions

5 lessons

Lesson 3 of 5

Oxidising and Reducing Agents

30 min

Two cleaners that must never meet

Household bleach contains sodium chlorate(I), NaClO. Many toilet cleaners contain hydrochloric acid. Pour one onto the other and a choking yellow-green gas rises from the bowl: chlorine. $\mathrm{ClO^-(aq) + Cl^-(aq) + 2H^+(aq) \rightarrow Cl_2(g) + H_2O(l)}$ Chlorine in the +1 state (in $\mathrm{ClO^-}$ ) takes an electron from chlorine in the −1 state (in $\mathrm{Cl^-}$ ), and both end at 0. The chlorate(I) ion acts as an oxidising agent, the chloride ion as a reducing agent. This lesson names the common oxidising and reducing agents, ranks their strength, and shows the colour changes that reveal them in the laboratory.

Definition

Oxidising agent (oxidant): a substance that oxidises another by taking electrons from it; the oxidising agent is itself reduced. Example: $\mathrm{Cl_2}$ in $\mathrm{Cl_2 + 2I^- \rightarrow 2Cl^- + I_2}$ .

Reducing agent (reductant): a substance that reduces another by giving it electrons; the reducing agent is itself oxidised. Example: $\mathrm{I^-}$ in the same reaction.

Redox titration: a titration in which an oxidising agent of known concentration reacts with a reducing agent (or the reverse), the end point being shown by a colour change.

Common oxidising agents

Oxidising agent	Reduction half-equation	Colour change
acidified potassium manganate(VII)	$\mathrm{MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O}$	purple → colourless
acidified potassium dichromate(VI)	$\mathrm{Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O}$	orange → green
chlorine	$\mathrm{Cl_2 + 2e^- \rightarrow 2Cl^-}$	pale green → colourless
hydrogen peroxide (acid)	$\mathrm{H_2O_2 + 2H^+ + 2e^- \rightarrow 2H_2O}$	none
iron(III) ions	$\mathrm{Fe^{3+} + e^- \rightarrow Fe^{2+}}$	yellow-brown → pale green
concentrated nitric acid	$\mathrm{NO_3^- + 2H^+ + e^- \rightarrow NO_2 + H_2O}$	brown gas given off
oxygen	$\mathrm{O_2 + 4H^+ + 4e^- \rightarrow 2H_2O}$	none

Common reducing agents

Reducing agent	Oxidation half-equation	What you may see
reactive metals (Zn, Mg)	$\mathrm{Zn \rightarrow Zn^{2+} + 2e^-}$	metal dissolves
iodide ions	$\mathrm{2I^- \rightarrow I_2 + 2e^-}$	colourless → yellow-brown (blue-black with starch)
iron(II) ions	$\mathrm{Fe^{2+} \rightarrow Fe^{3+} + e^-}$	pale green → yellow-brown
sulfur dioxide, sulfite	$\mathrm{SO_2 + 2H_2O \rightarrow SO_4^{2-} + 4H^+ + 2e^-}$	none
hydrogen sulfide	$\mathrm{H_2S \rightarrow S + 2H^+ + 2e^-}$	yellow sulfur forms
carbon, carbon monoxide, hydrogen	e.g. $\mathrm{CO + H_2O \rightarrow CO_2 + 2H^+ + 2e^-}$	used to reduce metal oxides
ethanedioate (oxalate) ions	$\mathrm{C_2O_4^{2-} \rightarrow 2CO_2 + 2e^-}$	gas bubbles

Some substances can act either way. Hydrogen peroxide oxidises iodide ions, but it reduces manganate(VII), being itself oxidised to oxygen: $\mathrm{H_2O_2 \rightarrow O_2 + 2H^+ + 2e^-}$ . Sulfur dioxide usually reduces, but it oxidises hydrogen sulfide to sulfur.

How strong?

The standard electrode potential, $E^\circ$ , of each half-equation measures how strongly its oxidised form pulls in electrons. The more positive $E^\circ$ , the stronger the oxidising agent (Figure 16.3). A substance can oxidise the reduced form of any couple below it in the list: chlorine (+1.36 V) oxidises iodide (+0.54 V), but iodine cannot oxidise chloride.

Figure 16.3: Standard electrode potentials of common oxidising agents (values from Wikipedia's data page). Red: strong oxidising agents used in the laboratory.

Tests in the laboratory

For a reducing agent: add a few drops of acidified potassium manganate(VII). A reducing agent turns the purple solution colourless. Acidified potassium dichromate(VI) turns from orange to green.
For an oxidising agent: add acidified potassium iodide solution. An oxidising agent turns it yellow-brown as iodine forms; a drop of starch then turns it blue-black.

Crystals of potassium manganate(VII), KMnO₄, in a small jar: dark purple, almost black, with a metallic sheen. In acid solution it is one of the strongest common oxidising agents (E° = +1.51 V) and its purple colour disappears as it is reduced. Photo: W. Oelen, CC BY-SA 3.0, via Wikimedia Commons

Worked example 1: who oxidises whom?

For the bleach reaction $\mathrm{ClO^- + Cl^- + 2H^+ \rightarrow Cl_2 + H_2O}$ , write the two half-equations, and identify the oxidising and reducing agents.

Solution. Oxidation numbers of Cl: +1 in $\mathrm{ClO^-}$ , −1 in $\mathrm{Cl^-}$ , 0 in $\mathrm{Cl_2}$ . $\text{reduction: } \mathrm{2ClO^- + 4H^+ + 2e^- \rightarrow Cl_2 + 2H_2O} \qquad \text{oxidation: } \mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$ Adding and dividing by 2 gives the overall equation. $\mathrm{ClO^-}$ is reduced, so it is the oxidising agent; $\mathrm{Cl^-}$ is oxidised, so it is the reducing agent. Two oxidation states meet at zero.

Worked example 2: iron in an iron tablet

An iron tablet is dissolved in dilute sulfuric acid, and the iron(II) ions are titrated with 0.0200 mol dm⁻³ potassium manganate(VII). The end point (first permanent pink colour) is reached after 12.00 cm³. Calculate the mass of iron in the tablet. (Fe = 55.85)

Solution. $\mathrm{MnO_4^- + 8H^+ + 5Fe^{2+} \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O}$ : 1 mol of manganate(VII) reacts with 5 mol of iron(II). $n(\mathrm{MnO_4^-}) = 0.0200 \times \frac{12.00}{1000} = 2.40 \times 10^{-4}\ \mathrm{mol} \qquad n(\mathrm{Fe^{2+}}) = 5 \times 2.40 \times 10^{-4} = 1.20 \times 10^{-3}\ \mathrm{mol}$ $m(\mathrm{Fe}) = 1.20 \times 10^{-3} \times 55.85 = 0.0670\ \mathrm{g} = 67.0\ \mathrm{mg}$ No indicator is needed: manganate(VII) is its own indicator, because its purple colour persists only when all the iron(II) has been used up.

Warning

Common mistakes. (1) Saying "the oxidising agent is oxidised": it is the opposite; the oxidising agent is reduced. (2) Forgetting acid in manganate(VII) reactions: without enough $\mathrm{H^+}$ , brown $\mathrm{MnO_2}$ forms instead of $\mathrm{Mn^{2+}}$ . (3) Using hydrochloric acid to acidify manganate(VII) titrations: $\mathrm{MnO_4^-}$ oxidises $\mathrm{Cl^-}$ to chlorine, so too much titrant is used; dilute sulfuric acid is used instead. (4) Treating every colour change as a test for oxidation: check which substance gains or loses electrons.

Safety. Never mix bleach with acids or acidic cleaners. Potassium manganate(VII) and potassium dichromate(VI) are strong oxidising agents; dichromate(VI) is also toxic and a cause of cancer, so schools use small volumes of dilute solution, with gloves and goggles, and collect the waste. If chlorine is released, leave the room, ventilate it, and get anyone who is coughing into fresh air and to medical help.

Tip

The agent does to others what is done to itself in reverse. An oxidising agent oxidises others and is reduced; a reducing agent reduces others and is oxidised. Purple → colourless or orange → green means a reducing agent was there; colourless iodide → brown means an oxidising agent was there.

The A Plus Squad: Identifying Reducing and Oxidising Agents, an Experiment (6 min). Watch acidified potassium manganate(VII) lose its purple colour with a reducing agent, and potassium iodide turn brown with an oxidising agent.

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Practice

Define an oxidising agent and a reducing agent in terms of electrons.
In $\mathrm{Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2}$ , name the oxidising agent and the reducing agent.
Describe what you would see when sulfur dioxide is bubbled through acidified potassium dichromate(VI), and explain the change.
Describe a test to show that a solution contains an oxidising agent.
Use Figure 16.3 to predict whether bromine will oxidise (a) iodide ions and (b) chloride ions. Explain.
Write the half-equation for hydrogen peroxide acting as a reducing agent, and name a substance it reduces.
25.0 cm³ of an iron(II) solution needs 20.0 cm³ of 0.0200 mol dm⁻³ $\mathrm{KMnO_4}$ for complete oxidation. Calculate the concentration of iron(II) ions.
Explain why dilute sulfuric acid, not hydrochloric acid, is used to acidify manganate(VII) titrations.

Show answersHide answers

An oxidising agent accepts electrons (and is reduced); a reducing agent donates electrons (and is oxidised).
Oxidising agent: $\mathrm{Fe_2O_3}$ (Fe +3 → 0). Reducing agent: CO (C +2 → +4).
The orange solution turns green. $\mathrm{SO_2}$ is a reducing agent: it reduces $\mathrm{Cr_2O_7^{2-}}$ (Cr +6, orange) to $\mathrm{Cr^{3+}}$ (+3, green) and is oxidised to sulfate.
Add acidified potassium iodide: an oxidising agent turns the colourless solution yellow-brown (iodine); starch then gives a blue-black colour.
(a) Yes: $E^\circ(\mathrm{Br_2}) = +1.07\ \mathrm{V}$ is more positive than $E^\circ(\mathrm{I_2}) = +0.54\ \mathrm{V}$ , so bromine takes electrons from iodide. (b) No: chlorine (+1.36 V) is a stronger oxidising agent than bromine, so bromine cannot take electrons from chloride.
$\mathrm{H_2O_2 \rightarrow O_2 + 2H^+ + 2e^-}$ ; it reduces acidified manganate(VII) (purple to colourless).
$n(\mathrm{MnO_4^-}) = 0.0200 \times 0.0200 = 4.00 \times 10^{-4}\ \mathrm{mol}$ ; $n(\mathrm{Fe^{2+}}) = 5 \times 4.00 \times 10^{-4} = 2.00 \times 10^{-3}\ \mathrm{mol}$ ; concentration $= \frac{2.00 \times 10^{-3}}{0.0250} = 0.0800\ \mathrm{mol\,dm^{-3}}$ .
Manganate(VII) is strong enough to oxidise chloride ions to chlorine, so some titrant would be used up by the acid and the result would be too high. Sulfate ions are not oxidised.

Quick check

Acidified potassium manganate(VII) is added to a solution and turns from purple to colourless. What does this show?

Key points

An oxidising agent accepts electrons and is reduced; a reducing agent donates electrons and is oxidised.
Common oxidising agents: acidified $\mathrm{MnO_4^-}$ (purple → colourless), acidified $\mathrm{Cr_2O_7^{2-}}$ (orange → green), $\mathrm{Cl_2}$ , $\mathrm{H_2O_2}$ , $\mathrm{Fe^{3+}}$ . Common reducing agents: reactive metals, $\mathrm{I^-}$ , $\mathrm{Fe^{2+}}$ , $\mathrm{SO_2}$ , $\mathrm{H_2S}$ , C, CO, $\mathrm{H_2}$ .
The more positive $E^\circ$ , the stronger the oxidising agent.
In manganate(VII) titrations, 1 mol $\mathrm{MnO_4^-}$ reacts with 5 mol $\mathrm{Fe^{2+}}$ ; manganate(VII) is its own indicator.

Sources and credits

Image: Potassium permanganate crystals.jpg by W. Oelen, via Wikimedia Commons, licensed under CC BY-SA 3.0.
Electrode potentials from Standard electrode potential (data page), Wikipedia contributors, licensed under CC BY-SA 4.0.
Video: Identifying reducing and oxidizing agents - experiment, The A Plus Squad, embedded from YouTube.
The text, tables, chart, worked examples and Practice items were written and drawn by Ganzaa.
This lesson uses CC BY-SA material and is shared under CC BY-SA 4.0.

1Two cleaners that must never meet

2Common oxidising agents

3Common reducing agents

4How strong?

5Tests in the laboratory

Two cleaners that must never meet

Common oxidising agents

Common reducing agents

How strong?

Tests in the laboratory