Chapter 14

Electrochemistry

Basic Concepts

Chemical Reaction that involves the transfer of electrons. A Redox reaction. Loss of electrons – oxidation Gain of electrons – reduction

Oxidizing agent. A species that takes electrons.

Reducing agent. A species that gives electrons.

Basics

Na(s) + H+ -> Na+ + H2(g) Sodium is a reducing agent Hydrogen ion is the oxidizing agent.

Basics

We are donating and gaining electrons. If we could use these electrons perhaps we could do some useful work.

If we can make the electron travel in an electrical circuit then the amount of current can be measured.

Current is related to reaction rate or amount of reaction

Potential is related to free energy change of the reaction.

Electron Charge

q used to denote. Unit is Coulombs (C) Charge on a single electron is

1.602x10-19 C which will allow us to determine the charge on a mole of electrons.

1.602x10-19 C * 6.022x1023 mol-1) = 96490 C mol-1

This is called the Faraday Constant q = nF n is the number of moles

Current

Charge flowing through a circuit

One ampere, the charge of one coulomb per second flowing past a given point.

Electrodes

The interface between a solution and an electrical circuit. Can be actively involved or just serve as a source or sink for electons.

Electrical Potential

Work required when moving and electric charge from one point to another.

Electrical potential (E) is measured in Volts (V).

Work is a measure of energy, measured in joules (J).

Work = E * q Joules volts coulombs

Free Energy

Maximum amount of work that can be done on the surroundings is equal to the Gibbs free energy change.

then G = -work = -Eq Or G = -nFE

Ohm’s Law

Current is proportional to the potential and and inversely proportional to the resistance.

I = E/R

Electric Circuit

Power

Work done per unit time. Unit is the J/s which is know as the watt (W).

P = work/sec = Eq/sec = E(q/sec) = EI

P = EI = I2R = E2/R

Galvanic Cells

Spontaneous chemical reaction used to generate electricity.

An example might be

Voltmeter

A device to measure electrical potential. When electrons tend to flow into the negative terminal then a positive voltage is measured.

In this cell 2 AgCl (s) + 2 e- = 2 Ag + 2 Cl- (aq) Red Cd (s) + = Cd2+ + 2 e- Oxidation Cd (s) + 2 AgCl (s) = Cd2+ + 2 Cl- Net For this reaction we have a G of -150 kJ/mole per

mole of Cd oxidized.

Potential of this System

G = -150 kJ/mole then we have

E = - G/nF = -150 x 103 J / (2 mol)(9.649x104 C/mol) E = + 0.777 J/C = +0.777 V

Cathode/anode

Cathode electrode where reduction occurs

Anode electrode where oxidation occur

Put both terms in alphabetical order to remember

Salt Bridge

Any bridge in upstate New York in the winter.

Used to isolate the half cells so the work can be forced out into an external circuit.

The following cell has a problem.

What is it?

The silver ions in solution can go directly to the cadmium electrode surface and be reduced there.

We need to put in a barrier to rapid ionic transfer.

What about this cell

Isn’t this cute

Chemistry paper dolls?

Line Notation - Instead of Having to Draw the Cells

| phase boundary || salt bridge For First Cell

Cd(s) | CdCl2(aq) | AgCl(s) | Ag(s)

For Second Cell Cd(s) | Cd(NO3)2(aq) || AgNO3(aq) | Ag(s)

A Word of Connectors (Two common in USA)

Standard Potential Eo The energy to a half cell at standard

conditions (1 M and 25 C)

Let us look at the reduction of silver ion. Ag+ + e- = Ag(s)

We will compare this to a fixed reference.

That is the SHE or NHE Standard or Normal Hydrogen Electrode. H+ (aq, A=1) + e- = ½ H2 (g, A = 1)

SHE - All other redox couples are compared to this half cell. It is assigned a value of 0.000 V

In our cell the left side electrode (Pt) is attached to the negative terminal. (Reference)

Value of E are collected into Tables (Appendix H)

Nernst Equation

For the half reaction aA + ne- = bB

aA

bBo

A

A

nF

RTEE ln

Eo = is the standard Potential

R = gas constant (8.314472 (V*C)/(k*mol)

T = Temp (K)

N = # of electrons in the half reaction

F = Faraday

A = Activity

We will often lump the constants and assume 25 C

Nernst equation (25 C and converting to log10

aA

bBo

A

A

n

VEE log

05916.0

Complete Reaction

E = E+ - E- for full cell Steps

Write both half cells as reductions, make electrons equal

Half cell connected to positive terminal is E+

Other half cell is E-

Net voltage is from the above equation Balance equation (reversing the left half reaction

and adding to other half cell) E > 0 spontaneous as written E < 0 spontaneous in reverse

Eo and K

Kn

E o log05916.0

05916.010

onE

K

Cells as Chemical Probes

Equilibria between the half cells

Equilibria within each half cell

A Probe Cell

Probe Cell

Right side: We have our Ksp equilibrium The electrochemical reaction under this is

AgCl(s) + e- = Ag(s) + Cl- (aq, 0.10 M) Eo = 0.222 v

Left side: We have our Ka for the weak acid. The electrochemical reaction

2 H+(aq) + e- = H2 (g, 1.00 bar) E = 0.00, but H+ is not fixed at 1 M so E varies with

H+

Eo’

Formal Potential Since so many redox couples exist in

the body and many have H+ we modify the potential that we use to pH 7. (A little more reasonable than 1 M acid.

Homework

14- 4 13, 14, 15 and 27