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 - PowerPoint PPT PresentationTRANSCRIPT
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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
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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
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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:
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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
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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
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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
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Fundamentals of Electrochemistry
Galvanic Cells
3.) Electrodes
Cathode: electrode where reduction takes place
Anode: electrode where oxidation takes place
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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
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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
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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
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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
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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
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Fundamentals of Electrochemistry
Standard Potentials
3.) Example: Calculate Eo for the following reaction:
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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
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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)
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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:
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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
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Fundamentals of Electrochemistry
Eo and the Equilibrium Constant
2.) Example: Calculate the equilibrium constant (K) for the following reaction:
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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.
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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:
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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
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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)