RedoxYear 12 Chemistry

What is Redox?• REDOX stands for REDuction/Oxidation• Which species are oxidised?

2Mg(s) + O2(g) 2MgO(s)Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

• Zn metal has been oxidised as it has lost electrons, and the Cu2+ has been reduced as it has gained electrons.

• Zn Zn2+ + 2e and Cu2++ 2e Cu

• Oxidation refers to a loss of electrons• Reduction refers to a gain of electrons

Redox reactions involve the transfer of electrons

Definitions for Redox ReactionsDefinitions for Redox Reactions

• OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen.

• REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.

• OXIDISING AGENT—accepts electrons (gets reduced) to facilitate oxidation of another species

• REDUCING AGENT—donates electrons (gets oxidised) to facilitate reduction of another species.

• OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen.

• REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.

• OXIDISING AGENT—accepts electrons (gets reduced) to facilitate oxidation of another species

• REDUCING AGENT—donates electrons (gets oxidised) to facilitate reduction of another species.

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• Oxidation and reduction always occur together

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Zn Zn2+ + 2e and Cu2++ 2e Cu

• Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.

Oxidant and ReductantZn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) Zn Zn2+ + 2e and Cu2+ + 2e Cu

• The species that is oxidised, in this case Zn, is called the reducing agent or reductant.

• The species that is reduced, in this case Cu2+, is the oxidising agent or oxidant.

• Which is reduced /oxidised and which is the reductant /oxidant?

• reduced oxidised• Fe3O4 + 4C 3Fe + 4CO• oxidant reductant

Oxidation Number RulesThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferred.

1. Free elements and molecules have an oxidation number of zero.

Na, Be, K, Pb, H2, O2, HCl, H2O = 0

2. In monatomic ions, the oxidation number is equal to the charge on the ion.

Li+ = +1; Fe3+ = +3; O2- = -2

3. The oxidation number of oxygen is usually –2. (In H2O2 and O2

2- it is –1.)

4. The oxidation number of hydrogen is +1 except when it is bonded to metals in hydrides (e.g. LiH). In these cases, its oxidation number is –1.

6. The sum of the oxidation numbers of all the atoms in a molecule or polyatomic ion is equal to the charge.

5. Fluorine is always –1.

HCO3-

O = -2 H = +1

3x(-2) + 1 + ? = -1

C = +4

Oxidation numbers of all the atoms in HCO3

- and S in H2SO4? S = +6

Remember charge on molecules = 0

2Mg (s) + O2 (g) 2MgO (s)0 0 2+ 2-

Recognising Oxidation and Reduction in a Redox reaction

1. Write the ON for each atom

2. Increase in ON means oxidation

3. Decrease in ON means reduction

S + O2 SO2

PCl3 + Cl2 PCl5

Zn(s) + 2HCl(aq) ZnCl2(s) + H2(g)

NaCl + AgNO3 NaNO3 + AgCl

Explain which species is oxidised? Use ON to support your answer.

ON of S increases from 0 to +4 therefore it is oxidised.

Is it redox?

• 2HCl + Ca(OH)2 CaCl2 + 2H2O

• Cr2O72- + 2OH- 2CrO4

2- + H2O

• Mg + Cl2 MgCl2

• Pb2+ + 2I- PbI2

• 2NH4NO3 2N2 + O2 + 4H2O

Use ON to prove whether these reactions are redox

Mg is oxidised as its ON increases from 0 in Mg to +2 in Mg2+

Cl- (not Chlorine) is reduced as its ON decreases from 0 in Cl2 to -1 in Cl-

Test YourselfQ- Define oxidation and reduction and represent

each as a chemical equation.A- oxidation = loss of e– … X X+ + e–

reduction = gain of e– … X + e– X– Q- Why are 2Na + Cl2 2NaCl & 2H2 + O2 2H2O

considered redox reactions?A- Both involve the transfer of electrons (Na,

Cl2 ,H2 and O2 have ON=0. After reaction ON are Na+ = 1, Cl- = -1, H+ = +1 and O2- = -2

Q- Is it possible to oxidise a material without reducing something else?

A- No. A lost e– is taken up by something else.

Test YourselfQ- Define oxidising and reducing agent.A- An oxidising agent causes oxidation by being

reduced itself and a reducing agent causes reduction by being oxidised itself.

Q- Explain using equations why Ca + Cl2 CaCl2 is a redox reaction.

A- CaCl2 is an ionic compound made of positive calcium ion and negative chlorine ionsCa Ca2+ + 2e–, Cl2 + 2e– 2Cl–. Thus Ca is losing electrons (oxidation) and Cl is gaining electrons (reduction).

Balancing Half Equations

Mg Mg2+ + 2e-

Cl2 + 2e- 2Cl-Oxidation half-reaction (lose e-)Reduction half-reaction (gain e-)

1. Write half equation by identifying reactant and product

2. Balance atoms that are not O or H

3. Balance O by adding H2O and H by adding H+

4. Balance charge by adding e- to the most positive side

MnO4- Mn2+

2I- I2

I- I2

MnO4- Mn2+

MnO4- + 8H+ Mn2+ + 4H2O2I- I2

2I- I2 + 2e- MnO4- + 8H+ +5e- Mn2+ + 4H2O

Balance the following, are they oxidation or reduction?

Cl2 Cl-

Balancing Redox Equations

1. Write the unbalanced equation for the reaction in ionic form.

The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?

Fe2+ + Cr2O72- Fe3+ + Cr3+

2. Separate the equation into two half-reactions.

Oxidation:

Cr2O72- Cr3+

+6 +3

Reduction:

Fe2+ Fe3++2 +3

3. Balance the atoms other than O and H in each half-reaction.

Cr2O72- 2Cr3+

Balancing Redox Equations4. For reactions in acid, add H2O to balance O atoms and H+ to

balance H atoms. Cr2O7

2- 2Cr3+ + 7H2O

14H+ + Cr2O72- 2Cr3+ + 7H2O

5. Add electrons to the most positive side of each half-reaction to balance the charges on the half-reaction.

Fe2+ Fe3+ + 1e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

6. If necessary, equalise the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.

6Fe2+ 6Fe3+ + 6e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

Balancing Redox Equations

7. Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. You should also cancel like species.

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

6Fe2+ 6Fe3+ + 6e-Oxidation:

Reduction:

14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O

8. Combine H+ and OH- to make water.

Balance the following Fe2+ +MnO4- Fe3+ +Mn2+

2H2O + 2e- H2 + 2OH-

2H20 O2 + 4H+ + 4e-

Halogens as Oxidants• Fluorine is so powerful an oxidant that it

oxidises water to oxygen.

2F2 + 2H2O 4HF + O2

A halogen higher in the Group can oxidise the ions of one lower down.

• Chlorine reacts with bromide and iodide

Cl2 + 2Br- 2Cl- + Br2

Cl2 + 2I- 2Cl- + I2

• Bromine reacts with iodide

Br2 + 2I- 2Br- + I2

• Iodine does not react with either of the other halide ions.

Test Yourself

• Write a balanced equation for Chlorine reacting with Magnesium and use oxidation numbers to explain which element is oxidised.

Cl2 + Mg Mg2+ + 2Cl-

• Mg changes its oxidation number from 0 to +2, an increase in oxidation number means it is oxidised. Chlorine decreases its oxidation number from 0 to -1 so it is reduced.

0 0 +2 -1

Test Yourself• Write a balanced equation for Bromine reacting with

Potassium Iodide and use oxidation numbers to explain which element is the oxidant.

Br2 + KI KBr + I2

• Bromine changes its oxidation number from 0 to -1, a decrease in oxidation number means it is reduced. Iodine increases its oxidation number from -1 to 0 so it is oxidised. An oxidant gets reduced to help something else become oxidised therefore Br2 is the oxidant.

0 +1-1 -1+1 0

Halogens can Oxidise Water• Halogens are not very soluble in water• They do react with water

Cl2(s) + H2O(l) HCl(aq) + HOCl(aq)

• Hypochlorous acid (HOCl) and hypochlorite (OCl-) are the main components of free active chlorine used in disinfectants and swimming pools.

Metals Reacting with Metal Ions

E.g. Magnesium in copper sulfate solution

Mg Mg2+ + 2e-

Cu2+ +2e- CuBut if the reaction doesn’t work water is getting reduced

E.g. Sodium in Zinc sulfate2Na 2Na+ + 2e-

2H20 + 2e- H2 + 2OH-

Metals will reduce metal ions if the metal is higher on the reactivity series than the ion.

Common Reductants

Reductant State Appearance Product State Appearance

Mg s Silver/grey Mg2+ aq colourless

C s black CO2 g colourless

CO g colourless CO2 g colourless

H2 g colourless H2O l colourless

Fe2+ aq Green Fe3+ aq orange

Br- aq colourless Br2 aq orange

I- aq colourless I2 s grey

SO2 g colourless SO42- aq colourless

Any metal will displace a less reactive one from solution

K>Na>Li>Ca>Mg>Al>Zn>Fe>Sn>Pb>Cu>Ag

Common Oxidants

Oxidant State Appearance Product State Appearance

O2 g colourless O2- aq colourless

I2 s grey I- aq colourless

Cl2 g green Cl- aq colourless

Fe3+ aq orange Fe2+ aq green

H2O2 l colourless H2O l colourless

Cr2O72- /H+ l orange Cr3+ l green

MnO4-/H+ l purple Mn2+ l colourless

Cl2 + Mg MgCl2

Cl2 + 2e- 2Cl- and Mg Mg2+ + 2e-

+++

–––

Electrolysis – used to separate ions• “Cells” are containers of liquid with electrodes:

• In “electrolytic cells”, electricity is used to force chemicals to undergo a redox reaction

Source of electricity

Molten or aqueous ions(Electrolyte)

Cell Electrode carbon or platinum

Anions

Oxidised

Anode is positive

Cations

Reduced

Cathode is negative

e-

The electrolytic cell

• Electric current forces charges on electrodes

+++

–––

Na+Na+

Cl– Cl–

• Na+ takes up an electron: Na+(l) + e– Na • Cl– gives up an electron: 2Cl–(l) Cl2 + 2e– • Electricity flows until ions are used up• Pure Na is deposited, Cl2 gas is produced

• Na+ is attracted to cathode, Cl– to anode

Reactivity Series• Whether you get the metal or

hydrogen during electrolysis depends on the position of the metal in the reactivity series:

E.g. copper chloride solution Anode chlorine Cathode coppersodium chloride solution Anode chlorine Cathode hydrogen

Hydrogen will be produced unless the metal is lower on the reactivity series.

Electrolysis of Water

Test Yourself

Ionic Solution At the Cathode At the Anode

CuCl2 Cu Cl2

CuSO4 Cu O2

NaCl H2 Cl2

H2SO4 H2 O2

HCl H2 Cl2