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Fundamentals of Electrochemistry

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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

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slide1
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
slide2
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

slide3
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:

slide4
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

slide5
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

slide6
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

slide7
Fundamentals of Electrochemistry
  • Galvanic Cells

3.) Electrodes

Cathode: electrode where reduction takes place

Anode: electrode where oxidation takes place

slide8
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

slide9
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

slide10
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

slide11
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

slide12
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

slide13
Fundamentals of Electrochemistry
  • Standard Potentials

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

slide14
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

slide15
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)

slide16
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:

slide17
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

slide18
Fundamentals of Electrochemistry
  • Eo and the Equilibrium Constant

2.) Example:

      • Calculate the equilibrium constant (K) for the following reaction:
slide19
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.

slide20
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)

slide21
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

slide22
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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