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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
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)