Fundamentals of Electrochemistry Introduction 1.)Electrical Measurements of Chemical Processes ...

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

Batteries produce 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 No free electrons 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)

Fnq Relation between charge and moles:

Coulombs molesemol

Coulombs

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

Net Reaction:

Reduction:

Oxidation:

AgCl(s) is reduced to Ag(s)Ag deposited on electrode and Cl-

goes into solution

Electrons travel from Cd electrode to Ag electrodeCd(s) is oxidized to Cd2+

Cd2+ goes into solution

Fundamentals of Electrochemistry

Galvanic Cells

2.) Cell Potentials Reaction is spontaneous if it does not require external energy

Reaction Type E Cell Type

Spontaneous + Galvanic

Nonspontaneous - Electrolytic

Equilibrium 0 Dead battery

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

ˆ Larger 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

Two half-cell reactions

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

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

Phase boundaryElectrode/solution interface

Solution in contact with anode & its concentration

Solution in contact with cathode & its concentration

2 liquid junctionsdue to salt bridge

cathode

Fundamentals of Electrochemistry

Galvanic Cells

5.) Short-Hand Notation Representation of Cells: by convention start with anode on left

Ag+ + e- Ag(s) Eo = +0.799V

Fundamentals of Electrochemistry

Standard Hydrogen Electrode (S.H.E)

Hydrogen gas is bubbled over a Pt electrode

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

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

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:

EEEcellWhere: 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 electrodeE- = electrode where oxidation occurs (anode)

Electrons always flow towards more positive potential

Place values on number line to determine the potential difference

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

aA + ne- bB Eo

For the given reaction:

The ½ cell reduction potential is given by:

a

bo

aA

bBo

]A[

]B[log

n

VEE

A

Aln

nF

RTEE

0.05916

Where: E = actual ½ cell reduction potential

Eo = standard ½ cell reduction potentialn = number of electrons in reactionT = 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”

d

bo

a

co

cell]D[

]B[log

n

.E

]A[

]C[log

n

.EEEE

059160059160

aA + ne- cC E+o

dD + ne- bB E-o

Consider the following ½ cell reactions:

Cell potential in terms of Nernst Equation is:

ba

dcoo

cell]B[]A[

]D[]C[log

n

.)EE(E

059160

Simplify:

ba

dco

cell]B[]A[

]D[]C[log

n

.EE

059160

Fundamentals of Electrochemistry

Eo and the Equilibrium Constant

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

Since Eo=E+o- E-

o:

At equilibrium Ecell =0:

Klogn

.Eo

059160

Definition of equilibrium constant

05916010 .nEo

K

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

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

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

Fundamentals of Electrochemistry

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

Cells as Chemical Probes

2.) Example: If the voltage for the following cell is 0.512V, find Ksp for Cu(IO3)2:

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 PathwaysMetabolic 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´

ba

dco

cell]B[]A[

]D[]C[log

n

.EE

059160

Need to express concentrations asfunction of Ka and [H+].

Cannot use formal concentrations!

Eo´ (V)