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Fundamentals of Electrochemistry

- Introduction
- 1.)Electrical Measurements of Chemical Processes
- Redox Reaction involves transfer of electrons from one species to another.
- Chemicals are separated
- Can monitor redox reaction when electrons flow through an electric current
- Electric current is proportional to rate of reaction
- Cell voltage is proportional to free-energy change
- Batteriesproduce a direct current by converting chemical energy to electrical energy.
- Common applications run the gamut from cars to ipods to laptops

Fundamentals of Electrochemistry

- Basic Concepts
- 1.)A Redox titration is an analytical technique based on the transfer of electrons between analyte and titrant
- Reduction-oxidation reaction
- A substance is reduced when it gains electrons from another substance
- gain of e- net decrease in charge of species
- Oxidizing agent (oxidant)
- A substance is oxidized when it loses electrons to another substance
- loss of e- net increase in charge of species
- Reducing agent (reductant)

(Reduction)

(Oxidation)

Oxidizing

Agent

Reducing

Agent

Fundamentals of Electrochemistry

- Basic Concepts
- 2.)The first two reactions are known as “1/2 cell reactions”
- Include electrons in their equation

3.) The net reaction is known as the total cell reaction

- Nofreeelectrons in its equation

4.) In order for a redox reaction to occur, both reduction of one compound and oxidation of another must take place simultaneously

- Total number of electrons is constant

½ cell reactions:

Net Reaction:

Fundamentals of Electrochemistry

- Basic Concepts

5.) Electric Charge (q)

- Measured in coulombs (C)
- Charge of a single electron is 1.602x10-19C
- Faraday constant (F) – 9.649x104C is the charge of a mole of electrons

6.) Electric current

- Quantity of charge flowing each second

through a circuit

- Ampere: unit of current (C/sec)

Relation between charge and moles:

Coulombs

moles

Fundamentals of Electrochemistry

- Galvanic Cells

1.) Galvanic or Voltaic cell

- Spontaneous chemical reaction to generate electricity
- One reagent oxidized the other reduced
- two reagents cannot be in contact
- Electrons flow from reducing agent to oxidizing agent
- Flow through external circuit to go from one reagent to the other

Reduction:

Oxidation:

Net Reaction:

AgCl(s) is reduced to Ag(s)

Ag deposited on electrode and Cl-

goes into solution

Cd(s) is oxidized to Cd2+

Cd2+ goes into solution

Electrons travel from Cd

electrode to Ag electrode

Fundamentals of Electrochemistry

- Galvanic Cells

2.) Cell Potentials

- Reaction is spontaneous if it does not require external energy

Potential of overall cell = measure of the tendency of a reaction to proceed to equilibrium

ˆLarger the potential, the further the reaction is from equilibrium and the greater the driving force that exists

Fundamentals of Electrochemistry

- Galvanic Cells

3.) Electrodes

Cathode: electrode where reduction takes place

Anode: electrode where oxidation takes place

Fundamentals of Electrochemistry

- Galvanic Cells

4.) Salt Bridge

- Connects & separates two half-cell reactions
- Prevents charge build-up and allows counter-ion migration

Salt Bridge

- Contains electrolytes not

involved in redox reaction.

- K+ (and Cd2+) moves to cathode with

e- through salt bridge (counter

balances –charge build-up

- NO3- moves to anode (counter

balances +charge build-up)

- Completes circuit

Two half-cell reactions

Fundamentals of Electrochemistry

- Galvanic Cells

5.) Short-Hand Notation

- Representation of Cells: by convention start with anode on left

Zn|ZnSO4(aZN2+ = 0.0100)||CuSO4(aCu2+ = 0.0100)|Cu

Phase boundary

Electrode/solution interface

anode

cathode

2 liquid junctions

due to salt bridge

Solution in contact with

anode & its concentration

Solution in contact with

cathode & its concentration

Fundamentals of Electrochemistry

- Standard Potentials

1.) Predict voltage observed when two half-cells are connected

- Standard reduction potential (Eo) the measured potential of a half-cell reduction reaction relative to a standard oxidation reaction
- Potential arbitrary set to 0 for standard electrode
- Potential of cell = Potential of ½ reaction
- Potentials measured at standard conditions
- All concentrations (or activities) = 1M
- 25oC, 1 atm pressure

Ag+ + e-»Ag(s) Eo = +0.799V

Standard Hydrogen Electrode (S.H.E)

Pt(s)|H2(g)(aH2 = 1)|H+(aq)(aH+ = 1)||

Hydrogen gas is bubbled over a Pt electrode

Fundamentals of Electrochemistry

- Standard Potentials

1.) Predict voltage observed when two half-cells are connected

As Eo increases, the more favorable the reaction and the more easily the compound is reduced (better oxidizing agent).

Reactions always written as reduction

Appendix H contains a more extensive list

Fundamentals of Electrochemistry

- Standard Potentials

2.) When combining two ½ cell reaction together to get a complete net reaction, the total cell potential (Ecell) is given by:

Where:E+ = the reduction potential for the ½ cell reaction at the positive electrode

E+ = electrode where reduction occurs (cathode)

E- = the reduction potential for the ½ cell reaction at the negative electrode

E- = electrode where oxidation occurs (anode)

Place values on number line to determine the potential difference

Electrons always flow towards more positive potential

Fundamentals of Electrochemistry

- Standard Potentials

3.) Example: Calculate Eo for the following reaction:

Fundamentals of Electrochemistry

- Nernst Equation

1.) Reduction Potential under Non-standard Conditions

- E determined using Nernst Equation
- Concentrations not-equal to 1M

For the given reaction:

aA + ne-»bB Eo

The ½ cell reduction potential is given by:

Where: E = actual ½ cell reduction potential

Eo = standard ½ cell reduction potential

n = number of electrons in reaction

T = temperature (K)

R = ideal gas law constant (8.314J/(K-mol)

F = Faraday’s constant (9.649x104 C/mol)

A = activity of A or B

at 25oC

Fundamentals of Electrochemistry

- Nernst Equation

2.) Example:

- Calculate the cell voltage if the concentration of NaF and KCl were each 0.10 M in the following cell:

Pb(s) | PbF2(s) | F- (aq) || Cl- (aq) | AgCl(s) | Ag(s)

Fundamentals of Electrochemistry

- Eo and the Equilibrium Constant

1.) A Galvanic Cell Produces Electricity because the Cell Reaction is NOT at Equilibrium

- Concentration in two cells change with current
- Concentration will continue to change until Equilibrium is reached
- E = 0V at equilibrium
- Battery is “dead”

Consider the following ½ cell reactions:

aA + ne-»cC E+o

dD + ne-»bB E-o

Cell potential in terms of Nernst Equation is:

Simplify:

Fundamentals of Electrochemistry

- Eo and the Equilibrium Constant

1.) A Galvanic Cell Produces Electricity because the Cell Reaction is NOT at Equilibrium

SinceEo=E+o- E-o:

At equilibrium Ecell =0:

Definition of

equilibrium constant

at 25oC

at 25oC

Fundamentals of Electrochemistry

- Eo and the Equilibrium Constant

2.) Example:

- Calculate the equilibrium constant (K) for the following reaction:

Fundamentals of Electrochemistry

- Cells as Chemical Probes

1.) Two Types of Equilibrium in Galvanic Cells

- Equilibrium between the two half-cells
- Equilibrium within each half-cell

If a Galvanic Cell has a nonzero voltage then the net cell reaction is not at equilibrium

Conversely, a chemical reaction within a ½ cell will reach and remain at equilibrium.

For a potential to exist, electrons must flow from one cell to the other which requires the reaction to proceed not at equilibrium.

Fundamentals of Electrochemistry

- Cells as Chemical Probes

2.) Example:

- If the voltage for the following cell is 0.512V, find Kspfor Cu(IO3)2:

Ni(s)|NiSO4(0.0025M)||KIO3(0.10 M)|Cu(IO3)2(s)|Cu(s)

Fundamentals of Electrochemistry

- Biochemists Use Eo´

1.) Redox Potentials Containing Acids or Bases are pH Dependent

- Standard potential all concentrations = 1 M
- pH=0 for [H+] = 1M

2.) pH Inside of a Plant or Animal Cell is ~ 7

- Standard potentials at pH =0 not appropriate for biological systems
- Reduction or oxidation strength may be reversed at pH 0 compared to pH 7

Metabolic Pathways

Fundamentals of Electrochemistry

- Biochemists Use Eo´

3.) Formal Potential

- Reduction potential that applies under a specified set of conditions
- Formal potential at pH 7 is Eo´

Eo´ (V)

Need to express concentrations as

function of Ka and [H+].

Cannot use formal concentrations!

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