Download Electrochemical Methods: Intro

Electrochemical Methods: Intro
•Electrochemistry Basics
•Electrochemical Cells
•The Nernst Equation
•Activity
•Reference Electrodes (S.H.E)
•Standard Potentials
•Note: I think you should read Chapter 14 of Harris FIRST, before
reading the Skoog book.
Review of the Basics
• Oxidation
• Reduction
– Loss of electrons
– Always occurs at the
anode
– Happens because of
the action of a
reducing agent
3
3
– Gain of electrons
(charge is reduced)
– Always occurs at the
cathode
– Happens because of
the action of the
oxidizing agent
2
Fe  Ce  Fe  Ce
4
Electric Charge (in Coulombs) and Work
• Voltage represents electrical
• The charge in
potential (potential to do
coulombs (q) is equal
work)
to the number of
moles of electrons (n) • If some total charge in
coulombs (q) is moved
times the Faraday
through some electrical
Constant (F)
potential (E, in volts V) then
work is done!
q (coulombs)  n (moles) x F (Faraday Constant)
Coulombs
F  9.649E10
mole of e Work (joules)  E (volts) x q (coulombs)
Ohm’s Law and Power
• Ohm’s law relates
electrical resistance,
current and
potential!
• Power is the work done in
some unit time (e.g. joules
of work per second)
• The units of Power are
Watts (W)
• Ohm’s law and power are
related!
E(potential)  I (current) x R (resistanc e, in Ohms, )
Work (joules) E x q
Power (in Watts) 

second
s
Exq
q
E x E xI
s
s
Let’s Work Some Problems
• A 6.00V battery is connected across a
2.00 K resistor, how many electrons
flow through the circuit per second?
• How many joules of heat (heat is work)
are produced per electron?
• What voltage would the battery need to
be to deliver a power at 100.0 Watts?
Electrochemical Cells
• A complete cell
contains:
– anode
– cathode
– completed circuit (for
electrons to flow)
– a salt bridge
(usually!)
– an electrolyte solution
– chemical species that
undergo reaction.
• There are two basic
electrochemical cells:
– A GALVANIC cell
uses spontaneous
chemical reactions to
generate electricity
– In a “electrolytic”
cell, an electrical
potential is applied to
the cell to drive some
reaction.
What is happening at the
electrode(s) and how do we
describe the cell?
Anode half - rxn : Zn 0(s)  Zn 2 (aq)  2 e 2
Cathode half - rxn : Cu(aq)
 2 e -  Cu0(s)
Complete Cell reaction : Zn (s)  Cu2 (aq)  Zn 2 (aq)  Cu0 (s)
in shorthand, we use symbols!
a single vertical line marks the phase difference
a double vertical line marks the salt bridge
Anode on the left, cathode on the right
Including the counter - ions tells us something about the solutions
Zn (s) | ZnSO 4(aq) || CuSO4(aq) | Cu(s)
The Standard Hydrogen Electrode (SHE)
• The basis by which all other measurements are made.
• Assigned a potential of zero by definition!
• Not practical for regular use
Standard Potentials
• Standardized potentials (Eo), listed as reductions, for all
half-reactions
• Measured versus the S.H.E (0)
• Used in predicting the action in either a galvanic cell or
how much energy would be needed to force a specific
reaction in a non-spontaneous cell
• Assumes an activity of one for the species of interest
(usually a fair approximation) at a known temperature
in a cell with the S.H.E.
• Assumes that the cell of interest is connected to the (+)
terminal of the potentiometer (voltmeter) and the
S.H.E. is connected to the (-) terminal
Better Oxidizing Agents
in upper left hand corner.
Better Reducing
Agents in lower
Right hand corner
A cell is constructed (all activities equal 1), in which iron (III)
is oxidizes copper (II). Draw a picture of the cell, write the
shorthand diagram, and calculate the voltage of the cell from
Standard Potentials!
Cells and the Nernst Equation
• For any cell, the measured potential between the anode
and the cathode can be calculated:
Ecell  E   E where E  is the calculatedpotential
for the half - cell(electrode) connectedto the ()
terminalof the potentiometer
(the right cell and the right side of the potentiometer, usually)
___________________________________________________
E - is the potentialfor the half - cell(electrode)
connectedto the (-) terminalof the potentiometer
(usuallythe left side of the potentiometer and the left cell)
Using Ecell
•
•
•
•
Write the calculation as described earlier
Calculate E+ and ECalculate Ecell
Write a balanced cell reaction, by adding the
two half-reactions
– Write out the right cell half reaction
– Write out the left cell half reaction and reverse it
– Add the two reactions together to get a net, balanced
cell reaction.
• If, you use the conventions described here,then:
– If Ecell>0, the reaction is spontaneous to the right
– If Ecell<0, the reaction is spontaneous to the left
Nernst Equation
• Accounts for potentials of cells where the
reagents are not at an activity of 1
– Remember that standard potentials are at A=1
• Accounts for the number of electrons
transferred in a reaction, the temperature of
the reaction, LeChatelier’s Principle and a
variety of other factors
• Used to calculate E+ and E- under non-standard
conditions
– Most real cases!
Generichalf - rxn : aA  ne -  bB
b


RT
A
o
B
Nernst Equation: E  E 
x ln  a 
nF
 AA 
or
RT
o
E E 
x ln Q
nF
where Q is the reactionquotient (eq. expression)
and at 25 C
0.05916V
o
E E 
x log Q
n
o
Arsenic solid is reduced to arsine (AsH3(g)). Write the Nernst
Equation for the reaction. Find E when pH = 3.00 and the
partial pressure of arsine is 1.00 torr.
The Nernst Equation for Complete Reactions…..
• Setup a series of two Nernst equations
– One for E+
– One for E-
• Solve each Nernst equation to get either
E+ or E-
• Add together to get Ecell
A galvanic cell is assembled in which the left cell is the anode
and a cadmium metal electrode is oxidized to cadmium ion in
0.010 cadmium nitrate. In the right cell, the cathode, silver
ion is reduced to silver metal on a silver metal electrode in
0.50M silver nitrate.
1. Draw the cell (both a picture and a schematic diagram)
2. Write the half and net cell reactions
3. Calculate the net cell voltage
4. Indicate in which direction the cell is spontaneous
Using the Nernst Equation to Solve for a Cell as a Chemical
Probe!
• Best done by example. Solve for some
unknown quantity!
• The cell is described as follows. Solve for
the concentration of chloride if the
measured cell voltage is 0.485 V
Pt (s) | H2 (g,1.00atm) | H

II Cl (aq,?M) | AgCl(s) | Ag(s)
(aq,pH  3.6)
-
“Complications” in Cells
• A variety of complicating factors can
influence the ability of a cell to have the
ideal potential value
• Some are always present, some can be
corrected for or evaluated
• Their influence is usually less significant
than the potential or current imparted by
the two half-reactions
Junction Potentials
• Arise because of the differing
mobility of different ions in
solution
• A charge (potential) develops
when one ion moves to an
electrode more rapidly than the
other ion
• Counteracted by the use of
strong electrolytes in solution
and/or in the salt-bridge
– The strong electrolyte helps
overcome mobility-induced
differences in charge
between two cells
• Rarely significant in
electroanalytical
chemistry
Concentration Polarization
• Arises because of differing analyte concentrations in solution when a
reaction is initiated
• An electrical layer is created at the electrode surface
• This layer resists the flow of electrical charge unless fresh ions are brought
to the electrode
– Stirring
– Diffusion
Ohmic Potential (IR Drop)
• Some electrical energy must be used to get ions moving
in solution
• This energy is not registered on the potentiometer,
because it is used to impart kinetic energy to the ions
• So, in reality the calculation of Ecell is:
Ecell  E cathode  Eanode - Eohmic
Eohmic  IR