Fundamentals of Electrochemistry

It’s shocking!

Electroanalytical Chemistry: group of analytical methods based upon electrical properties of analytes when part of an electrochemical cell

Potentiometry involves the measurement of potential for quantitative analysis.

Electrolytic electrochemical phenomena involve the application of a potential or current to drive a chemical phenomenon, resulting in some measurable signal which may be used in an analytical determination

• The majority of chemical reactions can be classified as one of two kind of major reaction types.

• Acid/Base Reactions: proton transfer

• Oxidation/Reduction (Redox) Reactions : electron transfer

If you think about it!......

Basic Concepts

Redox reactions involve a species which is oxidized and another that is reduced.

3223 VFeVFe

In the above, Fe3+ is reduced to Fe2+ . It is the oxidizing agent.

Since G<0 for this reaction we can say that V3+ wants the extra electron less than Fe3+.

33 FeV

22 Ag2Cd Galvanic Cell

An AsideWhy won’t this cell work?

Ag+ will go to leftelectrode and ask fore from Cd(s) directly.

Will this cell work?

How badly do the electrons want to flow?

q = n x F

Moles

Coulomb Moles Coulombs

E/RIsec

q

I = current in ampsR = resistance in ohmsE = potential difference in Volts

-nFEG

Voltaic Cells Electrochemical cells that use an oxidation-reduction reaction to generate an electric current are known as galvanic or voltaic cells.

Voltaic Cells

The voltaic cell consist of the two reactions.

reductionoxidation

We can only measure E for the full reaction.We can only measure E for the full reaction.We would like to calculate E for the half reactions.Before doing this, we must recognize the E depends on concentrations.

e2ZnZn(s) 2

(g)He2H2 2-

(g)HZnZn(s)H2 22

+

Zn(s)e2Zn2

(g)He2H2 2-

(g)HZnZn(s)H2 22

Or equivalently we can write the reactions as follows

Voltaic Cells

1

1

1

)(2

2

gH

H

Zn

Since reactants and products are in their standard states, wecall the E for this cell the standard reduction potential (Eo).Here Eo = .76V.

e2ZnZn(s) 2

(g)He2H2 2-

(g)HZnZn(s)H2 22

+

We arbitrarily define the potential for, one half reaction, the second reaction above to be exactly 0V when reactants and products are in their standard states.

Since Eo for the cell is the sum of Eo’s for the two half reactions we see that Eo for the first half reaction is .76V.

Oxi

dizi

ng P

ower

In

crea

ses

Voltaic Cells

1

1

1

)(2

2

gH

H

Zn

This voltaic cell on the previous slide is fully described withthe following notation

Pt(s)|1)Α(g,H|1)Α(aq,H||1)Α(aq,Zn|Zn(s) 22

S.H.E.||1)Α(aq,Zn|Zn(s) 2

e2ZnZn(s) 2

(g)He2H2 2-

(g)HZnZn(s)H2 22

+

Voltaic cells can be described by a line notation based on the following conventions.

Single vertical line indicates change in state or phase. Within a half-cell, the reactants are listed before the products. Activities of aqueous solns are written in parentheses after the symbol for the ion or molecule. A double vertical line indicates a junction between half-cells. The line notation for the anode (oxidation) is written before the line notation for the cathode (reduction).

Line Notation For Voltaic CellsLine Notation For Voltaic Cells

Zn | Zn2+(1.0 M) || Cu2+(1.0 M) | Cuanode

(oxidation)cathode

(reduction)

Electrons flow from the anode to the cathode in a voltaic cell. (They flow from the electrode at which they are given off to the electrode at which they are consumed.) Reading from left to right, this line notation therefore corresponds to the direction in which electrons flow.

The Nernst EquationThe Nernst Equation

Zn(s)2eZn2

The Nernst equation relates the potential of a cell in its standard state to that of a cell not in its standard state..

We know from Le Chatelier’s principle that increasing the concentration of Zn2+ should drive the reaction to the right. In other words it should decrease the potential of the half cell. The Nernst equation allows us to calculate this increase for the above half reaction as

2Zn

Zn(s)o ln2F

RTE E

Zn(s)24e2Zn 2

2

Zn

2Zn(s)o

2

ln4F

RTE E

bB

cCo ln

nF

RTE E

At 25oC this equation simplifies to

bB

cCo log

n

.05916VE E

cCnebB The Nernst Eq. for the reaction is

The Nernst Equation For Complete Cell

EE E

Here E+ and E- are the potentials of the half cells connected to the positive and negative terminals of potentiometer respectively. Let’s consider an example.

Taken from http://chemed.chem.purdue.edu/genchem

Voltaic Cells

1

1

1

)(2

2

gH

H

Zn

- +

2

2

Zn

Zn(s)

2

H

)(H log2

.05916V.76- -log

2

.05916V E g

E+ and E- are potentials of half cells connected to positive and negative terminals of potentiometer respectively

EE E

2

H

Zn)(H 22log

2

.0591676. E

g

Zn(s)e2Zn2

(g)He2H2 2-

(g)HZnZn(s)H2 22

Calculating Equilibrium ConstantsCalculating Equilibrium Constants

3324 FeCeFeCe

34 CeeCe 23 FeeFe

Eo=1.700VEo=0.767V

made up of the following two half reactions

Since Eo is greater for cerium this reaction will be the reductionreaction. The standard potential for the galvanic cell would be

(0.767)- 1.700E-EE ooo

Calculating Equilibrium Constants Calculating Equilibrium Constants ContinuedContinued

3324 FeCeFeCe

In a galvanic cell we would have

]][Ce[Fe

]][Ce[Feln

nF

RT-E-EE-EE

42

3300

-

At equilibrium E=0 and

nF

GE-E

ooo This connection to free

energy is important

1610K

C25at Klog 05916.]][Ce[Fe

]][Ce[Feln

F

RTE-E o

42

33oo

Calculating Equilibrium Constants for Calculating Equilibrium Constants for Nonredox Reactions Nonredox Reactions

-23

23 COFe)(FeCO s

This is a Ksp problem. Not a redox problem. Nonetheless we can use electrochemistry to calculate Ksp by considering

sp-2

32oo Klog

2

.059]][CO[Felog

2

.059E-E

0.44VEFe(s)2Fe

0.756VECOFe(s)2)(FeCOo2

o-233

e

es

spKlog

059.

.316-2

11059./.3162sp 1010K

(at 25oC)

Electrochemistry SkillsElectrochemistry Skills

• Understand how voltaic cells work.• Be able to calculate standard reduction potentials for

voltaic cells, given the chemical reactions.• Be able to describe a voltaic cell using the line notation

and visa versa. Know which way electrons flow and where the anode and cathode are.

• Know how to work with the Nernst Eq. to include concentration dependencies and calculate equilibirum constants