Electrochemistry is a branchof chemistry that studies chemicalreactions which take place ina solution at the interface of anelectron conductor (the electrode)and an ionic conductor(the electrolyte)

Electrochemistry

Electrochemical Cell

Electrolytic

Cell

Oxidation – A process in which anelement attains a more positiveoxidation state

Na(s) Na+ + e-

Reduction – A process in which anelement attains a more negativeoxidation state

Cl2 + 2e- 2Cl-

Gain Electrons = Reduction

An old memory device to remember oxidation and reduction like this…

LEO says GER

Lose Electrons = Oxidation

Oxidizing agent

The substance that is reduced is the oxidizing agent

Reducing agent

The substance that is oxidized is the reducing agent

Galvanic (Electrochemical) Cells

It is anelectrochemicalcell that deriveselectrical energyfrom spontaneousredox reactiontaking place within a cell.

e-

e- e- e-

Zn - Cu Galvanic

Cell

Zn Zn2+ + 2e- E0= -0.76V

Cu2+ + 2e- Cu E0= +0.34V

From the table of reduction potentials:

Cu2+ + 2e- Cu E = +0.34V

The less positive, or more negative reduction potential becomes the oxidation…

Zn Zn2+ + 2e- E = -0.76V

Zn + Cu2+ Zn2+ + Cu E0 = + 1.10 V

Line Notation

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

An abbreviated representation of an electrochemical cell

Anodesolution

Anodematerial

Cathodesolution

Cathodematerial

| |||

Calculating G0 for a Cell

0 (2 )(96485 )(1.10 )coulombs Joules

G mol emol e Coulomb

G0 = -nFE0

n = moles of electrons in balanced redox equation

F = Faraday constant = 96,485 coulombs/mol e-

Zn + Cu2+ Zn2+ + Cu E0 = + 1.10 V

0 212267 212G Joules kJ

Nernst Equation

Standard potentials assume a concentration of 1 M. The Nernst equation allows us to calculate potential when the two cells are not of 1.0 M.

At 25 C (298 K) the Nernst Equation is simplified this way:

E=E0-RT ln(oxidation)nF (reduction)

E=E0-0.0591 log(oxidation)F (reduction)

Equilibrium Constants and Cell Potential

At equilibrium, forward and reverse reactions occur at equal rates, therefore:

1. The battery is “dead”2. The cell potential,E0 is zero volts

0 0.05910 log( )volts E K

n

Modifying the Nernst Equation (at 25 C):

Zn + Cu2+ Zn2+ + Cu E0 = + 1.10 V

Calculating an Equilibrium Constant from a Cell Potential

0.05910 1.10 log( )

2volts K

(1.10)(2)log( )

0.0591K

37.2 log( )K

37.2 3710 1.58 10K x

Conductance and Conductivity

NaCl

Conductance and conductivity

• Conductance of a solution is the inverse of its resistance. Resistance is measured in ohms, W.

Conductance is expressed as the reciprocal of ohms and is called the Siemens, S (1S = 1W-1)

• Resistance L

A

where l is the distance between the electrodes and A is the area of the electrodes

• Resistance = L where is resistivity

A

LA

Conductance and conductivity

• Conductivity, k, is the inverse of resistivity so,

• k = 1 = L = c’ = S cm-1

RA R

• Conductivity is dependent on the number ofions, so we introduce molar conductivity

m = k ,where C is the molar concentration

C

Concentration dependence of molar conductivities

• For strong electrolytes molar conductivity decreases slightly with increasing concentration

For weak electrolytes, molar conductivity is small at moderate concentrations but high at very low concentrations. This is because at low concentrations the proportion of dissociated molecules is high.

m0

• When the ions are far apart theirinteractions can be ignored

• Limiting molar conductivity is due tomigration of cations (+ve ions) in onedirection and anions (-ve ions) in the otherunder very dilute conditions

• m0 = ++

0 + --0

• +0 and -

0 are respectively the ionicconductivities of the cations and anions

Kohlrausch’s law gives the empirical relationship between molar conductivity and concentration for strong electrolytes

Kohlrausch’s law

Electrolytic Processes

A negative cell potential, (-E0)

A positive free energy change, (+G)

Electrolytic processes are NOT spontaneous. They have:

An electrolytic process is the use of electrolysisindustrially to refine metals or compounds at a high purity and low cost.

Electrolysis of Water

2 22 4 4H O O H e

2 24 4 2 4H O e H OH

In acidic solution

Anode reaction:

Cathode reaction:

-1.23 V

-0.83 V

-2.06 V2 2 22 2H O H O

Electroplating of Silver

Anode reaction:

Ag Ag+ + e-

1. Solution of the plating metal

3. Cathode with the object to be plated2. Anode made of the plating metal

4. Source of current

Cathode reaction:

Ag+ + e- Ag

Electroplating requirements:

Faraday's laws of electrolysis arequantitative relationships based on theelectrochemical researches publishedby Michael Faraday in 1834.

Faraday’s 1st Law Of ElectrolysisThe amount of chemical reaction whichoccurs at any electrode duringelectrolysis by a current is proportionalto the quantity of electricity passedthrough the electrolyte.

Amount of electricity in coulombs required to reduce or oxidized 1 mole of substance

= n x F x 1mole

Faraday’s 2nd Law Of Electrolysis

The amounts of different substanceslibrated by the same quantity ofelectricity passing through the electrolyticsolution are proportional to their chemicalequivalent weight .

Amount deposit (gm) = m x Q

F x nm = atomic mass of substanceQ = Total electric charge passed through substance F = Faraday’s Constantn = electrons transferred per ion

(1) Determination of equivalent masses of elements

(2) Electron metallurgy(3) Manufacture of non-metals(4) Electro-refining of metals