Year 11 Chemistry ~ Unit 2

Oxidation and ReductionMany of the chemical reactions that play a

significant role in maintaining our environment are oxidation-reduction reactions or redox reactions.

Examples of redox reactions are the reactions that corrode metals, combustion reactions and photosynthesis and respiration.

Redox ReactionsOriginally, oxidation was described as a reaction

with oxygen.

When oxygen reacts with a substance, the substance is said to be oxidised.

A transfer of oxygen is referred to as redox reaction.

The substance that has gained oxygen has oxidised.

The substance that has lost oxygen has reduced.

Redox ReactionsOxidation and reduction occur

simultaneously during a redox reaction.Consider the redox reaction:

Fe2O

3(s) + 3CO(g) 2Fe(s) + 3CO

2(g)

Fe2O

3(s) + 3CO(g) 2Fe(s) + 3CO

2(g)

Oxidation – gain of oxygen

Reduction – loss of oxygen

Electron TransferNot all redox reactions involve the transfer

of oxygen.A more accurate description of a redox

reaction is the transfer of electrons.Oxidation is the loss of electrons.Reduction is the gain of electrons.

OIL RIG

Oxidation Is Loss Reduction Is Gain

Half EquationsConsider the Redox reaction:

)()(2)( 22 sgs MgOOMg

The magnesium will undergo oxidation by losing electrons to form magnesium ions:

eMgMg 22

22 24 OeO

The oxygen gas will undergo reduction by gaining electrons to form oxide ions:

Writing Overall Redox EquationsWhen writing equations for redox reactions

we usually write the half equations first.The number of electrons must be balanced

so that both half equations contain the same amount of electrons.

All the reactants from both half equations are written on the left hand side of the equation while all the products are written on the right.

The electrons are cancelled out.

Figure 16.4 Copper wire placed in a silver nitrate solution forms deposits of silver crystals.

Oxidants and ReductantsAn oxidant (or oxidising agent) is a

substance that causes another to be oxidised, and is itself reduced.

A reductant (or reducing agent) is a substance that causes another to be reduced, and is itself oxidised.

Quiz

1. Balance the following equations

a) Cr2O72-

(aq) + S2O32-

(aq) Cr 3+(aq) + S2O6

2-(aq)

b) Mn3+ (aq) MnO2 (s) + Mn2+

aq)

c) C3H8(aq) + O2(aq) CO2(aq) + H2O(l)

2. For the above equations, state what is the reducing agent and what is the oxidizing agent.

Classifying Redox ReactionsIt is relatively easy to identify whether

half equations are either oxidation or reduction reactions but sometimes it is not so easy to identify full equations as redox reactions.

Redox reactions can be identified by observing a change in Oxidation Numbers of the substances involved.

Oxidation Number Rules1. Free elements have an oxidation number equal to 0. Eg

Na(s), C(s), Cl2(g).2. In ionic compounds the oxidation number is equal to the

charge on the ion. Eg CaCl2: Ca2+=+2 and Cl- = -1.3. Oxygen usually has an oxidation number of -2 (except in

peroxides such as H2O2 where it is -1).4. Hydrogen has +1 (except in metal hydrides eg NaH

where it is -1). 5. The sum of oxidation numbers in a neutral compound is 0

and in a polyatomic ion is equal to the charge of the ion.

Determining whether a species has undergone reduction or oxidation can be done using oxidation numbers:

Oxidation = Increase in Ox #. Reduction = Decrease in Ox #.

3. Give the oxidation states of the following substancesa) Mn3+

b) MnO2

c) C3H8 d) O2

e) CO2 f) H2O(l)

g) MnO4- h) S2O3

2-

i) Mn 2+ j) S2O62-

(aq)

Using Oxidation Numbers to Identify Redox Reactions

Determining whether a species has undergone reduction or oxidation can be done using oxidation numbers:

Oxidation = Increase in Ox #. Reduction = Decrease in Ox #.

4. Write the 1/2 equations for Redox reactions inside a lead acid battery. Then state which reaction occurs at the cathode and anode.

5. Using oxidation states work out what is being oxidized and what is being reduced

a) 6Na (s)+ N2 (g) 2Na3N (s)

b) Mg (s)+ Cl2 (g) MgCl2 (s)

c) 4Fe (s)+ 3O2 (g) 2Fe2O3 (s)

d) Ca (s)+ C (s) CaC2 (s)

e) MnO4-(aq) + S2O3

2-(aq) Mn 2+

(aq) + S2O62-

(aq)

f) VO2+

(aq) + Zn (s) VO 2+(aq) + Zn2+

(aq)

ExampleFor the reaction:

)(2)(2)( 22 ggg COOCO +2 -2 0 +4 -2

Carbon in the carbon monoxide has gone from +2 to +4 which means that oxidation has occurred.

Oxygen in the oxygen gas has gone from 0 to -2 which means that reduction has occurred.

As both oxidation and reduction have occurred the reaction is a redox reaction.

Writing Half Equations Although most half equations are quite easy

to write, some involving polyatomic ions can be more difficult.

The following steps will make balancing these half equations easier:

1. Balance all elements except O and H in the half equation.

2. Balance the O atoms by adding water. 3. Balance the H atoms by adding H+ ions. 4. Balance the charge by adding electrons (e) and

then add states.

ExampleA green solution containing Fe2+ ions is mixed with a purple solution containing MnO4

- ions. Fe3+ and Mn2+ ions are formed.

Write a balanced equation for this reaction.

The half equation involving the iron ions is quite simple:

Fe2+(aq) Fe3+

+ e-

ExampleThe half equation involving the manganese is a little more difficult:

Step 1: Balance all elements except for O and HMnO4

- Mn2+

Step 2: Balance O by adding waterMnO4

- Mn2+ + 4H2O

Step 3: Balance H atom by adding H+

MnO4- + 8H+ Mn2+ + 4H2O

Step 4: Balance the charge with electronsMnO4

- + 8H+ + 5e- Mn2+ + 4H2O

ExampleTo complete the full balanced equation we must balance the electrons in each half equation:

Fe2+ Fe3+ + e-

MnO4- + 8H+ + 5e- Mn2+ + 4H2O

5Fe2+ 5Fe3+ + 5e-

Combine the two equations:MnO4

- + 8H+ + 5e- + 5Fe2+ Mn2+ + 4H2O + 5Fe3+ + 5e-

Cancel out the electrons:MnO4

- + 8H+ + 5Fe2+ Mn2+ + 4H2O + 5Fe3+

(X5)

Evidence for Electron TransferWhen a redox reaction takes place, the results can

be visible but it is not always possible to see the transfer of electrons.

To show the transfer of electrons the half reactions must be separated and joined by an external circuit.

A galvanometer can be placed within the circuit to measure the flow of electrons.

Such an experiment is called a galvanic cell.

Figure 16.7 The apparatus used to demonstrate electron flow during oxidation–reduction reactions.

Galvanic CellsGalvanic cells consist of two half cells.Each half cell must consist of an electrode t= conduct the

electrons and an electrolyte in which ions are free to move through the solution.

Oxidation will occur in one half cell and reduction will occur in the other.

The electrode at which oxidation occurs is called the anode.The electrode at which reduction occurs is called the cathode.The half cells are also connected by a salt bridge.A salt bridge contains an ionic compound that allows ions to

flow between the solutions to complete the circuit and prevent an over accumulation of charge.

Anions flow into the anode to balance out the positive charge formed from oxidation.

Cations flow into the cathode to balance out the negative charge formed from reduction.

The Electrochemical Series – Predicting Redox Reactions

Chemists have constructed a table of half equations for redox reactions that can be formed in half cells in the order of their reactability or their ability to Oxidise. This is called the Electrochemical Series.

The elements that reduce most readily are at the top of the series.

The elements that are lower in the series are more likely to oxidise.

The electrochemical series is only valid for the conditions from which it was formed (standard conditions).

The series can be used to predict what will happen when two specific half cells are combined to form a galvanic cell.

The half cell that is higher in the series will reduce and the one lower will oxidise.

Predicting Redox ReactionsThe series can be used to predict what will

happen when two specific half cells are combined to form a galvanic cell.

The half cell that is higher in the series will reduce and the one lower will oxidise.