Transcript Electrochemistry

Electrochemistry Review
Goals:
•In this lab you will accomplish three things: Build a standard
voltaic (galvanic) cell and measure the potential.
•Determine the concentration of an unknown solution using the
potential measured from a voltaic cell.
•Measure the voltage of a series of solutions with different
hydronium ion concentrations and thus learn something about the
relationship between the Nernst equation and the values obtained
from a pH meter.
Introduction
• Electrochemistry deals with chemical reactions that
produce electricity and chemical reactions that are
propagated by an electrical current.
• In a voltaic cell electricity is produced from the
chemical reaction
• In an electrolytic cell electricity is pumped into the cell
to get the chemical reaction to proceed.
– used for electroplating metals such as
• chrome, copper or silver onto other surfaces
• as well as for the isolation of materials.
Introduction
• The following equation shows the reaction for
the isolation of sodium metal and chlorine gas
from molten sodium chloride.
– To force the reaction to the right, current must be
applied to this electrolytic cell.
2 NaCl (s)
2 Na o (s) + Cl2 (g)
Introduction
• In the experiment we will do in the laboratory, we will deal with
only a voltaic cell.
• For every electrochemical reaction something must be oxidized
and something must be reduced.
• The electrode where oxidation occurs is called the anode
• The electrode where reduction occurs in called the cathode
• Whether the anode is positive or negative depends upon
whether the cell is electrolytic or voltaic.
• Since most of us are exposed to voltaic cells such as batteries in
our everyday lives, it is common to think of the anode as
negative but that is true ONLY in voltaic cells
• Figure 1 is a
schematic of a
voltaic cell.
•In this voltaic cell the zinc metal is
with the zinc ions and the copper
metal is with the copper ions with a
salt bridge connecting the two.
•The salt bridge allows for the
slow transfer of ions through a
medium such as agar to which an
electrolyte such as potassium
nitrate has been added.
•Note the two half-cells.
•Each half of the cell is isolated
from the other and the circuit is
completed by a salt bridge
connecting the two half cells.
•This keeps each half-cell isolated;
in other words, the reduction
reaction separated from the
oxidation reaction.
Line Notation:
• For voltaic cells, it is common to use a shorthand
notation to indicate what is going on in the cell.
• Below is given the shorthand for the voltaic cell in
Figure 1.
Zn | Zn2+ (1.0M) || Cu 2+ (1.0M) | Cu
Standard hydrogen electrode (SHE)
• By isolating a reaction in a half-cell it is possible to calculate what is
called a standard reduction potential.
• The standard hydrogen electrode (SHE) is comprised of a platinum
electrode immersed in 1.0 M [H+] solution with hydrogen gas at 1
atmosphere of pressure bubbled over the electrode.
• The SHE can act as an anode where H2 is oxidized to H+, or it can act
as a cathode where the H+ is reduced to H2.
• In either reaction the SHE is assigned a potential of ZERO volts.
• Standard reduction potentials are measured against the SHE with
the corresponding electrode in a solution that is 1.0M.
Table 1 shows some standard reduction potentials (SRP)
Element
Electrode Reaction
Volts
Mg
Mg2+ + 2 e-  Mg
-2.37
Al
Al3+ + 3 e-  Al
-1.66
Zn
Zn2+ + 2 e-  Zn
-0.763
Fe
Fe2+ + 2 e-  Fe
-0.44
Ni
Ni2+ + 2 e-  Ni
-0.25
Sn
Sn2+ + 2 e-  Sn
-0.14
Pb
Pb2+ + 2 e-  Pb
-0.126
H2
2H+ + 2 e-  H2
0.000
Cu
Cu2+ + 2 e-  Cu
+ 0.337
I2
I2 + 2 e-  2 I -
+0.535
Fe3+
Fe 3+ + 1 e-  Fe 2+
+0.771
Ag
Ag+ + 1 e-  Ag
+0.799
Cl-
Cl2 + 2 e-  2 Cl -
+1.36
• As the reduction potential voltages
increase so does the tendency for the
ease of reduction of the species.
• Thus the reactions at the bottom of the
table with high reduction potentials are
strong oxidizing agents.
What if the solutions are at some other
concentration or pressure?
• SRPs given for reactions are at [1.0 M] for
aqueous solutions and 1 atm for gases
• The Nernst equation (Equation 1) can be used
to calculate SRPs at different concentrations
and pressures.
Nernst equation
2.303  RT
EE 
log 10 Q
nF
0
• E is the observed potential
under non-standard conditions
•E0 is the standard potential
•the value 2.303 converts the
equation from natural logarithm
(ln) to log10
•R is 8.314 J/mol*K
•T is temperature in Kelvin
•n is the number of electrons
transferred
•F is Faraday’s constant
•Q is the reaction quotient.
•it is the ratio of the
multiplicative product of the
concentrations of the products
to the multiplicative product of
the concentrations of the
reactants, each raised to the
exponent of its stoichiometric
coefficient.
Nernst equation
• Assuming a temperature of 298K, Equation 1
simplifies to Equation 2.
0.0592
EE 
log 10 Q
n
0
Example Calculation:
• Consider as an example a cell formed from the two half
reactions Fe3+ (1.00M)/ Fe2+ (0.20M) and Cl2 (5 atm)/Cl- (2.0
x 10-2M).
• First the standard cell potential, E0, can be calculated using
the standard reduction potentials given in Table 1.
• A quick look at the table lets us know that for the standard
potential of the cell to be positive (a requirement for a
spontaneous reaction), the chlorine will be reduced to
chloride ions and the iron II will be oxidized to iron III.
• Therefore the equation involving Fe2+ and Fe3+ must be
reversed as well as the sign of the potential given in the
table.
Example Calculation:
• The two half reactions can be summed as below to determine E0.
• The factor of 2 must be used in the second equation so that no
electrons are left in the net ionic equation, but the potential is not
multiplied by the factor of 2 because potential is the ratio of energy
to charge so any factor is canceled.
Cl2 + 2e-  2Cl+1.36 V
2Fe2+  2Fe3+ + 2e-0.771 V
______________________________________
Cl2 + 2Fe2+  2Cl- + 2Fe3+
+0.589 V
What is Q?
• For this reaction,
Q = [Cl-]2 [Fe3+]2
[Cl2] [Fe2+]2
Calculating the Cell Potential
• Using Equation 2 and substituting in the
conditions for the nonstandard cell given in
the example above, the potential for the cell
can be calculated:
E = 0.589 - 0.0592 log10 [2.0 x 10-2]2 [1.00]2 = 0.669 V
2
[5] [ 0.20]2
Procedure I
• Construct a simple voltaic cell using 0.1 M copper
sulfate CuSO4 solution and 0.1 M zinc sulfate
ZnSO4 solution.
• You will not be using 1.0 M standard solutions as
you might have expected from the discussion of
SRP. The reason for this is that problems are
encountered with the experiment at
concentrations of this magnitude.
– You will use the Nernst equation to calculate the
standard potential of 0.1 M cells.
Procedure I
• Using the 0.1 M copper solution and a zinc solution of
unknown concentration, construct another voltaic cell
and record the potential as above.
– Use the Nernst equation to determine the concentration
of the zinc solution.
• Using the 0.1 M zinc solution and a copper solution of
unknown concentration, construct another voltaic cell
and record the potential.
– Use the Nernst equation to determine the concentration
of the copper solution.
Procedure II
• Calibrate a pH meter using a two point calibration and two
selected buffers (pH = 4 & pH = 7).
• Once you have calibrated the pH meter, switch the readings
from pH to millivolts and record the millivolts for the series
of buffers provided, including the pH = 4 and pH = 7.
• After you have collected the data, plot the voltage of each
solution in volts as a function of log10[H3O+].
• Refer to the section on Concentration Cells at the end of
section 21.4 in Silberberg if you need any help with the
final question for this lab.