Transcript Chapter 14: Introduction to Electrochemistry

Fundamentals of Electrochemistry
Introduction
1.) Electrical Measurements of Chemical Processes

Redox Reaction involves transfer of electrons from one species to another.
-

Can monitor redox reaction when electrons flow through an electric current
-

Chemicals are separated
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”

3.)
Include electrons in their equation
The net reaction is known as the total cell reaction

No free electrons in its equation
½ cell reactions:
Net Reaction:
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
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
Relation between
charge and moles:
q  nF
Coulombs moles
Coulombs
mol e 
6.) Electric current

Quantity of charge flowing each second
through a circuit
Ampere: unit of current (C/sec)
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:
2+
AgCl(s)
totoAg(s)
Cd(s)isisreduced
oxidized
Cd
Electrons
travel from
Cd
Ag deposited
onto
electrode
and ClCd2+ goes
into
electrode
Ag solution
electrode
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
 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
Anode: electrode
where oxidation takes
place
Cathode: electrode
where reduction 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.
2+) moves
TwoCd
half-cell
reactions
 K+ (and
to cathode with
e through salt bridge (counter
balances –charge build-up
 NO3- moves to anode (counter
balances +charge build-up)
 Completes circuit
Fundamentals of Electrochemistry
Galvanic Cells
5.) Short-Hand Notation

Representation of Cells: by convention start with anode on left
Phase boundary
Electrode/solution interface
anode
Zn|ZnSO4(aZN2+ = 0.0100)||CuSO4(aCu2+ = 0.0100)|Cu
Solution in contact with
anode & its concentration
2 liquid junctions
due to salt bridge
cathode
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
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Ag+ + e-  Ag(s)

Potentials measured at standard conditions
All concentrations (or activities) = 1M
25oC, 1 atm pressure
-
Standard Hydrogen Electrode (S.H.E)
Pt(s)|H2(g)(aH = 1)|H+(aq)(aH+ = 1)||
2
Hydrogen gas is bubbled over a Pt electrode
Eo = +0.799V
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:
E cell  E   E 
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)
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
For the given reaction:
aA + ne-  bB
Eo
The ½ cell reduction potential is given by:
b
A
RT
E  Eo 
ln B
a
nF
AA
at 25oC
0.05916 V
[ B ]b
EE 
log
n
[ A]a
o
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
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
dD + ne-  bB
E+o
E-o
Cell potential in terms of Nernst Equation is:
E cell  E   E  
E o
0.05916
[ C ] c  o 0.05916
[ B ]b

log
 E 
log
a
d

n
n
[ A]
[
D
]

Simplify:
E cell  ( E o

E o
0.05916
[C ]c [ D ]d
)
log
n
[ A]a [ B ]b




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:
0.05916
[C ]c [ D ]d
E cell  E 
log
n
[ A]a [ B ]b
o
At equilibrium Ecell =0:
Definition of
equilibrium constant
Eo 
0.05916
log K
n
K  10
nE o
0.05916
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 Ksp for 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
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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´

0.05916
[C ]c [ D ]d
E cell  E 
log
n
[ A]a [ B ]b
o
Need to express concentrations as
function of Ka and [H+].
Cannot use formal concentrations!
Eo´ (V)