Revision: Describe Oxidation-Reduction Processes, 3credits

Definitions of Oxidation-ReductionLoss/Gain of electrons Increase/Decrease of oxidation

numberDetermining oxidation numbers

Remember:• A redox reaction is any reaction involving a transfer of

electrons.• In all redox reactions, oxidation and reduction happen at

the same time.• Oxidation is loss of electrons/ increase

in oxidation number.• Reduction is gain of electrons/decrease

in oxidation number.• Oxidising agents (oxidants) are

themselves reduced.• Reducing agents (reductants) are

themselves oxidised.

Rules for oxidation numbers1. Oxidation number for elements is zero.

N2, O2, O3, Cu, S8

2. Oxidation number of monoatomic ions is the same as their charge

Al3+, Zn2+, Cd2+, Ag+

0 0 0 0 0

+3 +2 +2 +1

3. In polyatomic ions (NO3-),

the sum of oxidation numbers equals the charge of the ion. In compounds (HNO3) the sum of the oxidation numbers equals zero.

4. Oxidation number of oxygen in most compounds is –2. Exceptions: H2O2, (peroxides) –1

5. Oxidation number of hydrogen is +1Exceptions: bonded to metals LiH

O22-

6. Fluorine is always –1.Other halogens are also –1, except when they are bonded to O, then they are positive.

Rules for oxidation numbers

Assign oxidation numbers to all of the elements:

Li2O Li = O =PF3 P = F =HNO3 H = N =

Cr2O72- Cr = O =

O =

+1 -2

MnO4- Mn = O =

-1+3+1 -2+5

-2-2+7

+6

Revision: Describe Oxidation-Reduction Processes, 3credits

Electrochemical cells: Their properties Electrode potentials defined as standard

electrode potentials, Eo (unit: Volts, V)

Cell diagrams: Use of the symbols “/” (phase boundary)

“,” (same phase)and “//” (salt bridge) Half cells Order of notation

Electrochemical definitions Electrochemical cell: a cell in which

oxidation and reduction occur, often in separate compartments

Half cell: a single electrode in an solution containing ions

Electrode: the conductor placed in cells that transfer charge between the external circuit and the electrolyte

Electrochemical definitions Anode: electrode where oxidation occurs

(negative electrode) Cathode: electrode where reduction

occurs (positive electrode) Electrolyte: substances in the salt bridge

(usually liquids) that transfer charge by moving ions

Electrolytic cell: a cell that uses a supply of electricity to bring about a non spontaneous chemical reaction (year 12)

Electrochemical definitions Electromotive force, EMF, or Eº

cell : the potential difference across a voltage source when no current is following

Standard reduction (electrode) potential Eº standard electrode potential measured in volts under standard conditions (25oC, 1molL-1, 1 atm), which indicates the ability of a species to gain electrons

Writing cell diagrams:Start on the left (oxidation)Zn/Zn2+//Cu2+/CuTake care when the electrode is not taking part in the reactionsTake care that you separate phases with the symbol /Take care that species in the same phase are separated by a comma

The salt bridge allows the movement of ions between the two half cells so that charges can be balanced. It completes the electric circuit.If the voltmeter (that restricts

the current flowing) is replaced by a wire, the reactions will take place more quickly. Here the copper electrode would

gain mass, the zinc electrode would loose mass.

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

Give the half cell diagramPt/H2/H+//

Revision: Describe Oxidation-Reduction Processes, 3credits

Electrochemical cells: Calculations related to Electrochemistry Spontaneity of oxidation-reduction

reactions Applications involving electrochemical

cells (details of particular cells, eg dry cells, will be provided as required)

Calculating emfThe emf of an electrochemical cell is calculated using the following formula:

Eocell= Eo(RHE) – Eo(LHE)

note: do not change the sign of the standard potentials

If the emf is positive:The electron flow is from left to right and the oxidation takes place in the left half cell.If the emf is negative:The electron flow is from right to left and the oxidation takes place in the right half cell.

Calculating emf

greater Eo: Strongest Oxidant, Reduction reaction

lower Eo: Strongest Reductant, Oxidation reaction

Predicting reactionsTo predict whether reactions happen spontaneously, the emf is calculated.Does Zn react with Fe3+?Zn is the possible loser of electronsZn/Zn2+//Fe3+, Fe2+/CEo

cell= 0.77 – (- 0.76) = 1.53 VThe emf is positive, therefore the electron flow is from left to right and oxidation takes place in the left half cell.That means that Zn reacts spontaneously with Fe3+.

Does Zn2+ react with Fe2+?Fe2+ is the possible loser of electronsC/ Fe2+, Fe3+ //Zn2+/ZnEo

cell= - 0.76 -0.77 = -1.53 VThe emf is negative, therefore the electron flow is from right to left and oxidation takes place in the right half cell.That means that Zn2+ does not react spontaneously with Fe2+.

Fe2+/Fe3+ = 0.77

Zn2+ /Zn = - 0.76

The Lead Acid Cell - a rechargeable battery

When the battery is charged, lead (II) ions (Pb2+) in lead sulfate arereduced to Pb and oxidised to lead (IV) ions (Pb4+) in lead oxide.

Observation: A build up of lead at the anode and a build up of PbO2 at the cathode.

2PbSO4 + 2H2O Pb + PbO2 + 2H2SO4

When the battery is discharged (providing energy to the car), the reaction is reversed and PbSO4 is produced. This will build up on cathode and anode. If any PbSO4 falls off the plate (which happens after long use), then it can not react and the battery needs replacing.

The Dry Cell - Lechlanche cell

You do not need to know the details about the Dry Cell as there are different types and in the exam a different example may be chosen.

A dry cell is very compact, so it may be difficult to identify cathode, anode and the half cell reactions.

The electrolyte used is a paste made up of alkaline or acidic salts.For the Lechlanche cell above

identify:Cathode:Anode:Oxidation:Reduction:

graphiteZinc case

Zn to Zn2+Mn4+ to Mn3+

Exam Questions:They may involve discussing the suitability of a redox pair for the

construction of a dry cell. The production of a gas indicates no

suitability!

Revision: Describe Oxidation-Reduction Processes, 3credits

Redox reactions: Appearance and state of common

oxidants and reductants Calculations involving mole ratios

(titrations)→ go through examples in your book

Reduced form Oxidised formCu brown solid Cu2+ blue aq

SO2gas SO4

2- aq

Mn2+ aq H+/MnO4- purple aq

H2O2liquid O2

gas

H2O liquid H2O2liquid

Cr3+ blue/green aq Cr2O72- orange aq

Fe2+ pale green aq Fe3+ orange aq

Cl- aq Cl2pale green gas

Br- aq Br2red/orange liquid

H2gas H+ aq

Reduced form Oxidised formMnO2

brown solid H2O/MnO4- purple aq

MnO42- green aq OH-/MnO4

- purple aq

I- aq I2 in I- = I3- brown aq

I2 in I- = I3- brown aq IO3

- aq

H2S gas S yellow/white solid

Pb2+ aq PbO2brown solid

NO2brown gas NO3

- aq

C2O42- aq CO2

gas

S2O32- aq S4O6

2- aq

Br2red/orange liquid BrO3

- aq

In previous exam papers, students struggled to achieve because: they did not read the question properly they were not able to show the direction of electron flow could not assign oxidation numbers did not give V as the unit for Eo

cell, or used the wrong sign could not write standard cell diagrams, forgot inert electrodes did not know that a salt bridge completes the circuit, allows ion flow did not know what happens when the voltmeter is replaced by a wire used the term “dissolve” incorrectly when referring to the decrease of

mass of an electrode could not identify the oxidised and reduced form did not know the colours of species did not identify strongest reductant/oxidant (quoted incorrectly a

redox pair)