Post on 23-Dec-2015
Electroanalytical Chemistry: group of analytical methods based upon electrical properties of analytes when part of an electrochemical cell
Potentiometry involves the measurement of potential for quantitative analysis.
Electrolytic electrochemical phenomena involve the application of a potential or current to drive a chemical phenomenon, resulting in some measurable signal which may be used in an analytical determination
• The majority of chemical reactions can be classified as one of two kind of major reaction types.
• Acid/Base Reactions: proton transfer
• Oxidation/Reduction (Redox) Reactions : electron transfer
If you think about it!......
Basic Concepts
Redox reactions involve a species which is oxidized and another that is reduced.
3223 VFeVFe
In the above, Fe3+ is reduced to Fe2+ . It is the oxidizing agent.
Since G<0 for this reaction we can say that V3+ wants the extra electron less than Fe3+.
33 FeV
How badly do the electrons want to flow?
q = n x F
Moles
Coulomb Moles Coulombs
E/RIsec
q
I = current in ampsR = resistance in ohmsE = potential difference in Volts
-nFEG
Voltaic Cells Electrochemical cells that use an oxidation-reduction reaction to generate an electric current are known as galvanic or voltaic cells.
The voltaic cell consist of the two reactions.
reductionoxidation
We can only measure E for the full reaction.We can only measure E for the full reaction.We would like to calculate E for the half reactions.Before doing this, we must recognize the E depends on concentrations.
e2ZnZn(s) 2
(g)He2H2 2-
(g)HZnZn(s)H2 22
+
Zn(s)e2Zn2
(g)He2H2 2-
(g)HZnZn(s)H2 22
Or equivalently we can write the reactions as follows
Since reactants and products are in their standard states, wecall the E for this cell the standard reduction potential (Eo).Here Eo = .76V.
e2ZnZn(s) 2
(g)He2H2 2-
(g)HZnZn(s)H2 22
+
We arbitrarily define the potential for, one half reaction, the second reaction above to be exactly 0V when reactants and products are in their standard states.
Since Eo for the cell is the sum of Eo’s for the two half reactions we see that Eo for the first half reaction is .76V.
This voltaic cell on the previous slide is fully described withthe following notation
Pt(s)|1)Α(g,H|1)Α(aq,H||1)Α(aq,Zn|Zn(s) 22
S.H.E.||1)Α(aq,Zn|Zn(s) 2
e2ZnZn(s) 2
(g)He2H2 2-
(g)HZnZn(s)H2 22
+
Voltaic cells can be described by a line notation based on the following conventions.
Single vertical line indicates change in state or phase. Within a half-cell, the reactants are listed before the products. Activities of aqueous solns are written in parentheses after the symbol for the ion or molecule. A double vertical line indicates a junction between half-cells. The line notation for the anode (oxidation) is written before the line notation for the cathode (reduction).
Line Notation For Voltaic CellsLine Notation For Voltaic Cells
Zn | Zn2+(1.0 M) || Cu2+(1.0 M) | Cuanode
(oxidation)cathode
(reduction)
Electrons flow from the anode to the cathode in a voltaic cell. (They flow from the electrode at which they are given off to the electrode at which they are consumed.) Reading from left to right, this line notation therefore corresponds to the direction in which electrons flow.
The Nernst EquationThe Nernst Equation
Zn(s)2eZn2
The Nernst equation relates the potential of a cell in its standard state to that of a cell not in its standard state..
We know from Le Chatelier’s principle that increasing the concentration of Zn2+ should drive the reaction to the right. In other words it should decrease the potential of the half cell. The Nernst equation allows us to calculate this increase for the above half reaction as
2Zn
Zn(s)o ln2F
RTE E
Zn(s)24e2Zn 2
2
Zn
2Zn(s)o
2
ln4F
RTE E
bB
cCo ln
nF
RTE E
At 25oC this equation simplifies to
bB
cCo log
n
.05916VE E
cCnebB The Nernst Eq. for the reaction is
The Nernst Equation For Complete Cell
EE E
Here E+ and E- are the potentials of the half cells connected to the positive and negative terminals of potentiometer respectively. Let’s consider an example.
2
2
Zn
Zn(s)
2
H
)(H log2
.05916V.76- -log
2
.05916V E g
E+ and E- are potentials of half cells connected to positive and negative terminals of potentiometer respectively
EE E
2
H
Zn)(H 22log
2
.0591676. E
g
Zn(s)e2Zn2
(g)He2H2 2-
(g)HZnZn(s)H2 22
Calculating Equilibrium ConstantsCalculating Equilibrium Constants
3324 FeCeFeCe
34 CeeCe 23 FeeFe
Eo=1.700VEo=0.767V
made up of the following two half reactions
Since Eo is greater for cerium this reaction will be the reductionreaction. The standard potential for the galvanic cell would be
(0.767)- 1.700E-EE ooo
Calculating Equilibrium Constants Calculating Equilibrium Constants ContinuedContinued
3324 FeCeFeCe
In a galvanic cell we would have
]][Ce[Fe
]][Ce[Feln
nF
RT-E-EE-EE
42
3300
-
At equilibrium E=0 and
nF
GE-E
ooo This connection to free
energy is important
1610K
C25at Klog 05916.]][Ce[Fe
]][Ce[Feln
F
RTE-E o
42
33oo
Calculating Equilibrium Constants for Calculating Equilibrium Constants for Nonredox Reactions Nonredox Reactions
-23
23 COFe)(FeCO s
This is a Ksp problem. Not a redox problem. Nonetheless we can use electrochemistry to calculate Ksp by considering
sp-2
32oo Klog
2
.059]][CO[Felog
2
.059E-E
0.44VEFe(s)2Fe
0.756VECOFe(s)2)(FeCOo2
o-233
e
es
spKlog
059.
.316-2
11059./.3162sp 1010K
(at 25oC)
Electrochemistry SkillsElectrochemistry Skills
• Understand how voltaic cells work.• Be able to calculate standard reduction potentials for
voltaic cells, given the chemical reactions.• Be able to describe a voltaic cell using the line notation
and visa versa. Know which way electrons flow and where the anode and cathode are.
• Know how to work with the Nernst Eq. to include concentration dependencies and calculate equilibirum constants