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7/29/2019 redox hand out

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4.1 Oxidation Reduction Reactions Acid-Base equilibrium, complexation equilibrium and

solubility equilibrium. These three equilibria have something in common, that is there is no change in oxidationstate of the species involved in the reaction.

In an oxidationreduction reaction, also known as a redoxreaction, electrons are not shared, but are transferredfrom one reactant to another. As a result of this electrontransfer, some of the elements involved in the reactionundergo a change in oxidation state. Those speciesundergoing an increase in their oxidation state areoxidized, while those undergoing a decrease in theiroxidation state are reduced.

For example, in the following redox reaction between Fe3+

and oxalic acid, H2C2O4, iron is reduced since its oxidationstate changes from +3 to +2. 1

2Fe3+(aq) + H2C2O4(aq) + 2H2O(l) 2Fe2+(aq)+

2CO2(g)+2H3O+(aq)

Oxalic acid, on the other hand, is oxidized since theoxidation state for carbon increases from +3 in

H2C2O4 to +4 in CO2.

Redox reactions, such as the one shown in theabove equation, can be divided into separate half-reactions that individually describe the oxidationand the reduction processes.

H2C2O4(aq) + 2H2O(l) 2CO2(g) + 2H3O+(aq) +

2e-

Fe3+ (aq) + e- Fe2+(aq) 2

It is important to remember, however, that

oxidation and reduction reactions always occur in

pairs. This relationship is formalized by the

convention of calling the species being oxidized a

reducing agent, because it provides the electrons for

the reduction half-reaction. Conversely, the species

being reduced is called an oxidizing agent. Thus, in

reaction given above example, Fe3+ is the oxidizing

agent and H2C2O4 is the reducing agent.3

Redox reaction: is an electron-transfer

reaction.

Oxidation: loss of electrons

Reduction: gain of electrons.

Reducing agent: A species that donates

electrons to another species.

Oxidizing agent: A species that accepts

electrons from another species4

An electrochemical cell is a device that permits

interconvertion of chemical energy into electrical

energy.

2.4.2 Electrochemical Cells

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Chemicalenergy Electrical

energy

Two kinds of Electrochemical cells

Galvanic cells( spontaneous)

Electrolytic cells(electrolysis)

( non-spontaneous

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Important terms about electrochemical cells(Galvanic Cell)

e

ee

e

ee

() (+)

Oxidation

half-cellReduction

half-cell

Pt or CInert or activeZn, Mg, etc.

Anions

Cations

Electrolytes

A salt bridge serves three purposesTo prevent direct reaction between R.A. and O.A.Tocomplete the circuit by allowing flow of ions

Anions

Cations

To balance the charge of ions in each half-ce ll

Loss of ve charge replenished by anionGain of -ve charge is cancelled by cations

Cations

Anions

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Standard Electrode Potentials

Because the voltage associated with a given reactiongives a measure ofthe reactions tendency to takeplace, it would be useful to have a set of voltage fordifferent half reactions. These could then becombined to give the voltage for large number of

redox reactions. Unfortunately, the voltage producedby a single electrode in a cell cannot be measureddirectly. Solution to the problem is to choose onereference electrode and arbitrarily assign somevoltage to it. Then voltages can be assigned to anyother electrode merely by measuring the voltageproduced by a cell utilizing the electrode in questionplus the reference electrode. The reference electrodechosen by international agreement is the standardhydrogen electrode. 8

The standard hydrogen electrode

A standard electrode is one in which all reactants

and products of the electrode half reaction are in

their standard states. The standard state for an ion

in a solution is the one for which the activity of the

ion is defined as being unity. This is the ion at

1mol/L concentration in an ideal solution

(activities approximated by concentrations). The9

standard hydrogen electrode is constructed with 1mol/L

HCl solution, hydrogen gas (H2) at 1atm pressure as show

in the diagram.

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H2(g) at 1 atmand 298 K

Platinum electrodecoated withplatinum black

Solution containing1M H+(aq) at 298K

Outletfor H2(g)

Figure 5.3 Schematic diagram of the standard hydrogen electrode (SHE).

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The potential or voltage arbitrarily assigned to the

standard hydrogen electrode is 0 V whether it operates

as an anode or cathode.

When a cell is constructed with standard hydrogen

electrode plus some other standard electrode (all

reactants and products in their standard state), the

measured potential is assigned solely to the otherelectrode.

Example:1. A standard copper-copper ion electrode is

combined with a standard hydrogen electrode to make

a cell. The cell voltage is measured as 0.34 V at 25C,

and electrons are found to enter the external circuit

from the hydrogen electrode.

a) Write the cell diagram for the cell

b) Write the overall redox reaction of the cell. 12


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Since the electrons leave the cell from the hydrogenelectrode, it must be the anode. The cell diagram andreactions are:

Pt(s)/H2(g)/H+(aq)//Cu2+/Cu(s)

Anode half reaction H2 (g) 2H+ (aq) + 2e-E = 0

Cathode half reaction Cu2+ (aq) + 2e- Cu (s)E = ?

H2 (g) + Cu2+

(aq) 2H+

(aq) + Cu (s) E = 0.34VEcell = Ecathode - EanodeEcathode = Ecell + Eanode = 0.34 + 0 = 0.34VSince an anode reaction is an oxidation, the potential

produced at such an electrode is called oxidationpotential. Similarly, the potential produced at a cathodeis called reduction potential. Either oxidation potential orreduction potentials could be assembled in a table, butby international agreement the latter are tabulated, asstandard reduction potentials.

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c) What is the potential for the standard copper-copper ion electrode?

After the standard potentials of any electrode

have been determined that electrode can be

used with another to find its potential. This ishow Table of Standard Reduction Potentials can

be complied.

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Table of Standard Reduction Potentials can be usedfor:

Predicting the voltage that a given standardgalvanic cell would produce.

Predicting the spontaneity of a given redoxreaction.

If the standard electrode potential is positive theredox reaction is spontaneous.

Other wise it is non-spontaneous.

Comparing the relative strengths of oxidizing andreducing agents.

Example: Zn can displace Cu in electrochemicalreaction

Li can displace Fe in electrochemical reaction15

Redox Equilibrium

Unlike the reactions that we have already considered,

the equilibrium position of a redox reaction is rarely

expressed by equilibrium constant. Since redox

reactions involve the transfer of electrons from a

reducing agent to an oxidizing agent, it is convenient to

consider the thermodynamics of the reaction in terms

of the electron.

The free energy, G, associated with moving a charge,

Q, under a potential, E, is given by

G = EQ 5.5.1

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Charge is proportional to the number of

electrons that must be moved. For a reaction in

which one mole of reactant is oxidized or

reduced, the charge, in coulombs, is

Q = nF 5.5.2

where n is the number of moles of electrons per

mole of reactant, and Fis Faradays constant

(96,485 C mol1 ). The change in free energy (in

joules per mole; J/mol) for a redox reaction,

therefore, is

G = nFE

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where G has units ofjoules per mole. Theappearance of a minus sign in equation is due to adifference in the conventions for assigning thefavored direction for reactions. In thermodynamics,reactions are favored when G is negative, andredox reactions are favored when Eis positive.

G =

G+RT ln Q 5.5.4

The relationship between electrochemical potentialand the concentrations of reactants and productscan be determined by substituting equation 5.5.3into the equation 5.5.4

nFE=nFE+RTln Q 5.5.5

where Eis the electrochemical potential understandard-state conditions. Dividing through out by nFleads to the well-known Nernst equation.

lnQnF

RTEE

o= 18


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Effect of Concentration on Cell Potential

G =G0 + RTlnQ

G0

= -nFE0

cell

-nFEcell= -nFE0

cell + RTln Q

Ecell= E0

cell - RTln Q

nF

Ecell= E0

cell - 0.0257ln Q

n

Ecell= E0

cell 0.0592log Q

n

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Nernst equation: An equation relating electrochemicalpotential to the concentrations of products andreactants.

The standard-state electrochemical potential, E,provides an alternative way of expressing the

equilibrium constant for a redox reaction. Since areaction at equilibrium has a G of zero, theelectrochemical potential, E, also must be zero.Substituting into equation 5.5.7 and rearranging showsthat

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Kn

Kn

ECell

log0592.0

ln0257.00

==

Example:

Calculate:

1. the standard-state potential,

2. equilibrium constant and

3. the potential when [Ag+] = 0.020 mol/L and

[Cd2+] = 0.050 mol/L, for the following reaction

taking place at 25 C.( EAg+/Ag = 0.7996 V,

ECd2+/Cd = -0.4030 V)

Cd(s) + 2Ag+ (aq) Cd2+ (aq) + 2Ag(s)

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Solution:

a) In this reaction from the given reactants, which one isoxidized? And which one is reduced?

Cd2+ is undergoing oxidation, and Ag+ is undergoingreduction. The standard-state cell potential, therefore, is

E = EAg+/Ag- ECd2+/Cd = 0.7996 V-(-0.4030 V) = 1.2026V

b) To calculate the equi librium constant, we substitute thevalues for the standard-state potential and number ofelectrons into equation .

Solving for Kgives the equilibrium constant as log K=40.6558

K = 4.527 x 1040

c) The potential when the [Ag+] is 0.020 mol/L and the [Cd2+]is 0.050 mol/L is calculated using Nernst equationemploying the appropriate relationship for the reactionquotient Q.

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