Redox Reactions

1. Define oxidation and reduction
Define the oxidation number from formulae
Describe tests for oxidising and reducing agents
Distinguish between oxidising and reducing agents
Chapter 11
Redox Reactions
LEARNING OUTCOMES

2.  For example, when magnesium is burned in oxygen, it
changes into magnesium oxide. We say that the
magnesium is oxidised into magnesium oxide.
 Oxidation can be defined as the gain of oxygen by a substance.
 The magnesium has gained oxygen to
become magnesium oxide.
2Mg(s) + O2(g)  2MgO(s)
Oxygen added
Magnesium + Oxygen  Magnesium oxide
Oxidation as the gain of oxygen
Chapter 11
Redox Reactions

3. For example, when copper(II) oxide is heated with
hydrogen, it changes to copper. We say that the
copper(II) oxide has been reduced to copper.
 Reduction can be defined as the loss or removal of
oxygen from a substance.
Reduction as the loss of oxygen
Chapter 11
Redox Reactions

4. Oxygen removed
 The copper(II) oxide has changed into
copper by its loss of oxygen.
CuO(s) + H2(g)  Cu(s) + H2O(l)
Copper(II) oxide + Hydrogen  Copper + water
Reduction as the loss of oxygen
Chapter 11
Redox Reactions

5. Hydrogen removed
Oxidation may also be defined as the loss or removal of
hydrogen from a substance.
H2S(g) + Cl2(g) S(s) + 2HCl(g)
 We say that the hydrogen sulphide is oxidised
to sulphur, because it has lost hydrogen.
 For example, hydrogen sulphide reacts with
chlorine to form sulphur and hydrogen chloride:
Oxidation as the loss of hydrogen
Chapter 11
Redox Reactions

6. Conversely, reduction may be defined as the gain or addition
of hydrogen to a substance.
N2(g) + 3H2(g)  2NH3(g)
Hydrogen added
 In this reaction, nitrogen is reduced to
ammonia, because it has gained hydrogen.
 For example, nitrogen reacts with hydrogen to
form ammonia in the Haber process:
Reduction as the gain of hydrogen
Chapter 11
Redox Reactions

7. Oxygen
added
In a redox reaction, if one substance is oxidised, the other is
being reduced.
E.g. The extraction of iron from iron(III) oxide in the blast
furnace:
Fe2O3(s) + 3CO(g) 2Fe(l) + 3CO2(g)
Fe2O3 loses oxygen,
and is thus reduced.
CO gains oxygen,
and is thus oxidised.
We say that iron(III) oxide is reduced to iron, and carbon
monoxide is oxidised to carbon dioxide.
Redox Reactions always occur together
Chapter 11
Redox Reactions

8. Hydrogen
added
For example, in the reaction of hydrogen sulphide with
chlorine:
H2S loses hydrogen,
and is thus oxidised.
Cl2 gains hydrogen,
and is thus reduced.
We say that hydrogen sulphide is oxidised to sulphur, and
chlorine is reduced to hydrogen chloride.
H2S(g) + Cl2(g)  S(s) + 2HCl(g)
Redox reactions always occur together
Chapter 11
Redox Reactions

9. Summary
Oxidation Reduction
Gain of oxygen Loss of oxygen
Loss of hydrogen Gain of hydrogen
Chapter 11
Redox Reactions

10. Quick check 1
1. State which substance is oxidised. What substance has it
oxidised to? Give a reason for your answer.
(a) C + O2  CO2
(b) Mg + H2O  MgO + H2
(c) 2CO + O2  2CO2
(d) H2I + Cl2  2HCl + I2
(e) CuO + H2  Cu + H2O
(f) Cl2(g) + H2S(g)  2HCl(g) + S(s)
(g) 2NH3 + 3CuO  3Cu + N2 + 3H2O Solution
Chapter 11
Redox Reactions

11. Quick check 1 (cont’d)
2. State which substance is reduced. What substance has it been
reduced to? Give a reason for your answer.
(a) ZnO + H2  Zn + H2O
(b) CO2 + 2Mg  2MgO + C
(c) Mg + H2O  MgO + H2
(d) Fe2O3 + 3CO  2Fe + 3CO2
(e) H2 + Cl2  2HCl
(f) CuO + Mg  Cu + MgO
(g) FeS + 2HCl  FeCl2 + H2S
Solution
Chapter 11
Redox Reactions

12.  We define:
 Oxidation is the loss of electrons from an
atom or ion.
 Reduction is the gain of electrons by an atom or ion.
 Redox reactions can take place even if
no oxygen or hydrogen is involved.
 A redox reaction is deemed to occur if there is
a transfer of electron(s) during the reaction.
Redox reactions in terms of electron transfer
Chapter 11
Redox Reactions

13.  For example, when sodium and chlorine react to form sodium
chloride:
 The sodium atom has transferred its outermost electron
to chlorine to form sodium chloride.
 The sodium atom has lost an electron, hence it is oxidised.
 The chlorine atom has gained an electron, hence it is reduced.
Chapter 11
Redox Reactions
Redox reactions in terms of electron transfer

14. 2Na + Cl2 2Na+
+ 2Cl-
Na loses electrons (oxidation)
Cl2 gains electrons (reduction)
 We say that sodium is oxidised (loss of electron) and chlorine is
reduced (gain of electron) to form sodium chloride.
Example 1: Reaction of sodium with chlorine
Chapter 11
Redox Reactions
Redox reactions in terms of electron transfer

15. Example 2: Reaction of magnesium with hydrochloric acid
Mg + 2H+
Cl-
 Mg2+
Cl-
2 + H2
H+
gains electrons (reduction)
 We say that magnesium is oxidised to
magnesium chloride. (loss of electrons)
 We say that hydrochloric acid is
reduced to hydrogen. (gain of electron).
Mg loses electrons (oxidation)
Chapter 11
Redox Reactions
Redox Reactions In Terms of Electron Transfer

16. 2Fe2+
Cl-
2 + Cl2  2Fe3+
Cl-
3
Example 3: Reaction of iron(II) chloride with chlorine.
Fe2+
loses electron to become Fe3+
(Oxidation)
Cl gains electron to become Cl-
(Reduction)
 Iron(II) chloride is oxidised to iron(III) chloride (loss of electrons)
 Chlorine is reduced to iron(III) chloride (gain of electrons)
Chapter 11
Redox Reactions
Redox reactions in terms of electron transfer

17.  To determine if an atom or ion has gained or lost electrons, we can
look at its oxidation state (or oxidation number).
 All free (uncombined) elements are assigned an oxidation state of
zero:
E.g. Na0
, Mg0
, Fe0
, Cu0
, H2
0
, Cl2
0
, O2
0
 The oxidation state of an element in a compound is equal to the
charge on the ion:
 E.g. H+
, Na+
, K+
(oxidation state +1);
Cl-
, Br-
, I-
(oxidation state -1);
Mg2+
, Ca2+
, Zn2+
, Fe2+
(oxidation state +2);
O2-
, S2-
, (oxidation state -2);
Fe3+
, Al3+
(oxidation state +3)
Oxidation States
Chapter 11
Redox Reactions

18.  When an atom or ion loses an electron, it is oxidised and its
oxidation state increases:
E.g. Na0
 Na+
+ e-
(From 0  +1)
E.g. Fe2+
 Fe3+
+ e-
(From +2  +3)
 When an atom or ion gains an electron, it is reduced and its
oxidation state decreases:
E.g. Cl0
+ e-
 Cl-
(From 0  -1)
E.g. Mg2+
+ 2e-
 Mg (From +2  0)
Redox reactions as changes in
oxidation state
Chapter 11
Redox Reactions

19. Example 1: Reaction of magnesium with hydrochloric acid
Step 1: Write down the balanced chemical equation.
Step 2: Write down the oxidation number of each
atom or ion in the equation.
Mg + 2H Cl  Mg Cl2 + H2
0 + - 2+ - 0
Redox reactions as changes in
oxidation state
Chapter 11
Redox Reactions

20. Step 3: Look for an atom or ion which has changed its oxidation
number in going from left to right in the equation.
Mg + 2H Cl  Mg Cl2 + H2
0 + - 2+ - 0
Step 4: Determine whether it is oxidation (increase in
oxidation state) or reduction (decrease in
oxidation state).
Oxidation (from 0 to +2)
Reduction (from + 1 to 0)
Redox reactions as changes in
oxidation state
Chapter 11
Redox Reactions

21. 2K+
I−
+ Cl2
0
 2K+
Cl−
+ I2
0
Potassium iodide isPotassium iodide is oxidisedoxidised to iodine.to iodine.
(( increaseincrease in oxidation state)in oxidation state)
Chlorine isChlorine is reducedreduced to KClto KCl
(( decreasedecrease in oxidation state)in oxidation state)
Example 2: Reaction of potassium iodide with chlorine.
 Notice that there is no change in K+
(in KI) to K+
(in KCl);
hence the potassium ion has not been oxidised or reduced.
Redox reactions as changes in
oxidation state
Chapter 11
Redox Reactions

22. KMnO4
KK++
(+1)(+1) xx
4(O4(O2-2-
))
(-2)(-2)
 Atoms in covalent and complex compounds can be given
oxidation states, assuming they are ionic.
 Oxidation states of all atoms in a compound must add up to zero
 Example: Find the oxidation state of Mn in KMnO4.
+1 + x + 4(-2) = 0
x = +7
Determination of Oxidation States
in a Compound
Chapter 11
Redox Reactions

23. Oxidation Reduction
Gain of oxygen Loss of oxygen
Loss of hydrogen Gain of hydrogen
Loss of electron(s)
(Increase in oxidation state)
Gain of electron(s)
(Decrease in oxidation state)
Summary
Chapter 11
Redox Reactions

24. 1. State which substance is oxidised. What substance has it been
oxidised to? State a reason for your answer.
(a) Zn + 2HCl  ZnCl2 + H2
(b) Mg + H2SO4  MgSO4 + H2
(c) Fe + Cl2  FeCl2
(d) Zn + CuSO4  ZnSO4 + Cu
(e) Fe + Pb(NO3)2  Fe(NO3)2 + Pb
(f) 2KI + Br2  2KBr + I2
Solution
Quick check 2
Chapter 11
Redox Reactions

25. 2. State which substance is reduced. What substance has it
been reduced to? State a reason for your answer.
(a) CuO + Mg  MgO + Cu
(b) 2Fe3+
+ 2Cl-
 2Fe2+
+ Cl2
(c) 2Na + Cl2  2NaCl
(d) Zn + CuSO4  ZnSO4 + Cu
(e) Mg + H2SO4  MgSO4 + H2
3. State the oxidation state of nitrogen in the following:
(i) NO, (ii) N2O, (iii) NO2, (iv) NO3
-
Solution
Quick check 2 (cont’d)
Chapter 11
Redox Reactions

26. 2Mg(s) + O2(g) 2MgO(s)
 In the above reaction, magnesium is
oxidised into magnesium oxide by oxygen.
 Consider the burning of magnesium in
oxygen to form magnesium oxide:
 Oxygen is called the oxidising agent.
Oxidising Agents and Reducing Agents
Chapter 11
Redox Reactions

27. An oxidising agent is a substance which causes
oxidation. It acts as an acceptor of electrons.
2Mg(s) + O2(g) 2MgO(s)
 In the above reaction, oxygen has received or
accepted 2 electrons from magnesium to
form magnesium oxide.
 Hence oxygen is the oxidising agent.Hence oxygen is the oxidising agent.
Definition:
Oxidising Agents
Chapter 11
Redox Reactions

28.  Other examples of oxidising agents are:
 chlorine and bromine
 potassium manganate(VII)
 potassium dichromate(VI)
Oxidising Agents
Chapter 11
Redox Reactions

29.  Consider the reaction between heated
copper(II) oxide and hydrogen.
CuO(s) + H2(g) Cu(s) + H2O(g)
 Copper(II) oxide is reduced to copper by
hydrogen.
 Hydrogen is called the reducing agentreducing agent..
Reducing Agents
Chapter 11
Redox Reactions

30. A reducing agent is a substance which causes
reduction. It acts as a donor of electrons.
CuO(s) + H2(g) Cu(s) + H2O(g)
 In the above reaction, hydrogen has given
away (donated) 2 electrons to the
copper(II) ion which then becomes copper.
 Hence hydrogen is the reducing agent.Hence hydrogen is the reducing agent.
Definition:
Reducing Agents
Chapter 11
Redox Reactions

31.  Other examples of reducing agents are:
 carbon
 carbon monoxide
 reactive metals like potassium, sodium, magnesium and aluminium
 potassium iodide
Reducing Agents
Chapter 11
Redox Reactions

32.  Since redox reactions always occur together, an oxidising agent will
be the substance reduced in the reaction.
 Similarly, a reducing agent will be the substance oxidised in the
reaction.
H2S(g) + Cl2(g)  S(s) + 2HCl(g)
HH22S is oxidised toS is oxidised to
sulphur bysulphur by chlorine.chlorine.
Chlorine is reduced to HClChlorine is reduced to HCl
byby hydrogen sulphide.hydrogen sulphide.
HH22S is the reducing agent.S is the reducing agent.ClCl22 is the oxidising agent.is the oxidising agent.
Oxidising Agents and Reducing Agents
Chapter 11
Redox Reactions

33. (a) Which substance is oxidised?
Ans: ________________________________________
(b) Which substance is reduced?
Ans: ________________________________________
(c) Which is the oxidising agent?
Ans: ________________________________________
(d) Which is the reducing agent?
Ans: ________________________________________
Worked Example
Fe2O3(s) + 3CO(g) 2Fe(l) + 3CO2(g)
Carbon monoxide is oxidised (gain of oxygen)
Iron(III) oxide is reduced (loss of oxygen)
Iron(III) oxide is the oxidising agent.
Carbon monoxide is the reducing agent.
Consider the
following reaction:
Chapter 11
Redox Reactions

34. Test for oxidising agent
 To test if an unknown substance is an oxidising agent, add a
solution of potassium iodide to it.
 If the mixture turns reddish brown due to the liberation of iodine
from the potassium iodide, then the unknown substance is an
oxidising agent.
Potassium iodide
solution added
unknown
solution
Mixture turns
reddish brown
Chapter 11
Redox Reactions

35.  To test if an unknown substance is a reducing agent, add an
acidified solution of potassium dichromate(VI)solution of potassium dichromate(VI) to it.
 If the mixture turns from yellow/orange to green,green, then the
unknown substance is a reducing agentreducing agent.
Test for reducing agent
Chapter 11
Redox Reactions

36. 1. In each of the following reactions, state
(i) the substance oxidised, (ii) the substance reduced, (iii) the
oxidising agent and (iv) the reducing agent.
(a) ZnO + CO  Zn + CO2
(b) Al2O3 + 3Mg  2Al + 3MgO
(c) 2FeCl2 + Cl2  2FeCl3
2. (a) Define oxidation in terms of electron transfer.
(b) Give an example of a redox reaction, including a
chemical equation with state symbols.
State clearly in your example, which substance is
oxidised and which substance is reduced.
Solution
Quick check 3
Chapter 11
Redox Reactions

37. Solution to Quick check 1
1. (a) C + O2  CO2
Carbon is oxidised into carbon dioxide. (gain of oxygen)
(b) Mg + H2O  MgO + H2
Magnesium is oxidised into magnesium oxide. (gain of oxygen)
(c) 2CO + O2  2CO2
Carbon monoxide is oxidised into carbon dioxide. (gain of oxygen)
(d) H2I + Cl2  2HCl + I2
Hydrogen iodide is oxidised into iodine. (loss of hydrogen)
(e) CuO + H2  Cu + H2O
Hydrogen is oxidised into water. (gain of oxygen)
(f) Cl2(g) + H2S(g)  2HCl(g) + S(s)
Hydrogen sulphide is oxidised into sulphur. (loss of hydrogen)
(g) 2NH3 + 3CuO  3Cu + N2 + 3H2O
Ammonia is oxidised into nitrogen. (loss of hydrogen)
Return
Chapter 11
Redox Reactions

38. 2.
(a) ZnO + H2  Zn + H2O
Zinc oxide is reduced into zinc. (loss of oxygen)
(b) CO2 + 2Mg  2MgO + C
Carbon dioxide is reduced into carbon. (loss of oxygen)
(c) Mg + H2O  MgO + H2
Water is reduced into hydrogen. (loss of oxygen)
(d) Fe2O3 + 3CO  2Fe + 3CO2
Iron(III) oxide is reduced into iron. (loss of oxygen)
(e) H2 + Cl2  2HCl
Chlorine is reduced into hydrogen chloride. (gain of hydrogen)
(f) CuO + Mg  Cu + MgO
Copper(II) oxide is reduced into copper.(loss of oxygen)
(g) FeS + 2HCl  FeCl2 + H2S
Iron(II) sulphide is reduced to hydrogen sulphide. (gain of hydrogen)
Return
Solution to Quick check 1 (cont’d)
Chapter 11
Redox Reactions

39. 1. (a) Zn + 2HCl  ZnCl2 + H2
Zinc is oxidised into zinc chloride.
(loss of electrons/increase in oxidation state)
(b) Mg + H2SO4  MgSO4 + H2
Magnesium is oxidised to magnesium sulphate.
(loss of electrons)
(c) Fe + Cl2  FeCl2
Iron is oxidised to iron(II) chloride. (loss of electrons)
(d) Zn + CuSO4  ZnSO4 + Cu
Zinc is oxidised to zinc sulphate. (loss of electrons)
(e) Fe + Pb(NO3)2  Fe(NO3)2 + Pb
Iron is oxidised to iron(II) nitrate. (loss of electrons)
(f) 2KI + Br2  2KBr + I2
Return
Solution to Quick check 2
Chapter 11
Redox Reactions

40. 2. (a) CuO + Mg  MgO + Cu
Copper(II) oxide is reduced to copper.
(loss of oxygen/decrease in oxidation state/gain of electrons)
(b) 2Fe3+
+ 2Cl-
 2Fe2+
+ Cl2
Iron(III) is reduced to iron(II). Decrease in oxidation state/gain of electron.
(c) 2Na + Cl2  2NaCl
Chlorine is reduced to sodium chloride. (gain of electron)
(d) Zn + CuSO4  ZnSO4 + Cu
Copper(II) sulphate is reduced to copper (gain of electrons)
(e) Mg + H2SO4  MgSO4 + H2
Sulphuric acid is reduced to hydrogen (gain of electron)
3. (i) +2, (ii) +1, (iii) +4, (iv) +5 Return
Solution to Quick check 2 (cont’d)
Chapter 11
Redox Reactions

41. 1. (a) ZnO + CO  Zn + CO2
(i) carbon monoxide, (ii) zinc oxide,
(iii) zinc oxide, (iv) carbon monoxide
(b) Al2O3 + 3Mg  2Al + 3MgO
(i) magnesium, (ii) aluminium oxide,
(iii) aluminium oxide, (iv) magnesium
(c) 2FeCl2 + Cl2  2FeCl3
(i) iron(II) chloride, (ii) chlorine,
(iii) chlorine, (iv) iron(II) chloride
2. (a) Oxidation occurs when there is a loss of electrons from an atom or ion.
(b) 2KI(aq) + Cl2 (g)  2KCl(aq) + I2(s)
Potassium iodide is oxidised to iodine.
Return
Solution to Quick check 3
Chapter 11
Redox Reactions

42. 1. http://library.kcc.hawaii.edu/external/chemistry/redox_title.html
2. http://www.chemistry.co.nz/redox_new.htm
3. http://en.wikipedia.org/wiki/Redox
To learn more about Redox reactions,
click on the links below!
Chapter 11
Redox Reactions