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CHEMISTRY 205

ELECTROCHEMISTRY:Voltaic Cell Measurements

I. Purpose:

• To get acquainted with the use of a voltmeter.

• To measure the cell potentials of different galvanic cells.

• To construct a concentration cell and measure its potential.

• To compare the observed cell potentials with the literature values and expected ones from Nernst equation.

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II. Theory:• Electrochemistry: the area of chemistry that

deals with the inter-conversion of electrical energy and chemical energy.

• Electrochemical processes: redox reactions in which the energy released by a spontaneousreaction is converted to electricity.

• Electrochemical cell: the experimental apparatus for generation of electricity through the use of a redox reaction.

• A galvanic cell, also called a voltaic cell, is an electrochemical cell in which electricity is produced by means of a spontaneous reaction.

A. Generation of Electricity:• A redox reaction occurs when the reducing

agent is in contact with the oxidizing agent.

• The electrons are transferred directly from the reducing agent to the oxidizing agent in solution.

• If these agents are physically separated from one another, the transfer of electrons can take place via an external conducting medium. As the reaction progresses, it sets up a constant flow of electrons and hence generates electricity.

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B. Example of a Galvanic Cell:• Consider a zinc bar dipped into a ZnSO4

solution, and a copper bar dipped into a CuSO4solution. The bars are called electrodes. By definition:anode: electrode at which oxidation occurs cathode: electrode at which reduction occurs

• Zn electrode, anode: Zn (s) � Zn2+(aq) + 2e-• Cu electrode, cathode: Cu2+(aq) + 2e- � Cu (s)

______________________________Zn (s) + Cu2+(aq) � Zn2+(aq) + Cu (s)

Electrochemical Cells

spontaneousredox reaction

anodeoxidation

cathodereduction

Ref. Chang Fig. 19.1

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• The two solutions must be separated from each other; otherwise, the Cu2+ ions will react directly with the zinc bar, and no useful electrical work will be obtained.

• However, to complete the electric circuit, the solutions must be connected by a conducting medium, a salt bridge, an inverted U tube containing an inert electrolyte such as KCl or NH4NO3 in agar.

• During the reaction, electrons flow externally from the anode (Zn electrode) through the wire and voltmeter to the cathode (Cu electrode).

• In the solution, the cations (Zn2+, Cu2+ and K+) move toward the cathode, while the anions (SO4

2-

and Cl-) move in the opposite direction, toward the anode.

• By convention:-The negative electrode in a galvanic cell is the electrode from which electrons are emitted (i.e. where oxidation occurs, anode). It is connected to the minus pole of the voltmeter.

-The positive electrode is the electrode to which the electrons are transferred (i.e. where reduction occurs, cathode). It is connected to the plus pole of the voltmeter.

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C. Electromotive Force:• The fact that electrons flow from one electrode

to the other indicates that there is a voltage difference between the two electrodes, called the electromotive force, or emf (E), which can be measured after connecting both electrodes to a voltmeter.

• The emf of a galvanic cell, also referred to as cell potential or cell voltage, is a direct measure of the driving force or thermodynamic tendency of the spontaneous redox reaction to occur.

• The cell potential depends on:-the nature of the electrodes and the ions, -the concentrations of the ions, -the temperature at which the cell is operated.

• The cell potential is the sum of the two electrode potentials. It is not possible, therefore, to measure individual absolute electrode potentials.

• However, a usual procedure is to assign a value of 0.00 Volts to the standard potential for the electrode reaction,

2H+(aq)+ 2e- � H2(g) EH+/,H2 = 0.000V @25°C

• Then, other electrode potentials can be given definite values based on the assigned value.

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D. Standard Electrode Potentials:Standard reduction potential (E0) is the voltage associated with a reduction reactionat an electrode when all solutes are 1 M and all gases are at 1 atm.

E0 = 0 V

Standard hydrogen electrode (SHE)

2e- + 2H+ (1 M) H2 (1 atm)

Reduction Reaction

Ref. Chang, Fig. 19.3

Ref. Chang, Fig. 19.4 (a)

E0 = 0.76 Vcell

Standard emf (E0 )cell

0.76 V = 0 - EZn /Zn0 2+

EZn /Zn = -0.76 V0 2+

Zn2+ (1 M) + 2e- Zn E0 = -0.76 V

E0 = EH /H - EZn /Zncell0 0

+ 2+2

Standard Electrode Potentials (continued):

E0 = Ecathode - Eanodecell0 0

Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

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Ref. Chang, section 19.3

• E0 is for the reaction as written

• The more positive E0 the greater the tendency for the substance to be reduced

• The half-cell reactions are reversible

• The sign of E0 changes when the reaction is reversed

• Changing the stoichiometriccoefficients of a half-cell reaction does not change the value of E0

E. Determination of Cell Potential:

• E°cell= E°red + E°ox

With E°M/M2+ = -E°M2+ /M

• E°cell= E°Cu2+/Cu + E°Zn/Zn2+

• E°cell= E°Cu2+/Cu - E°Zn2+/Zn

+0.34 –(- 0.76) = + 1.10V

E0 = Ecathode - Eanodecell0 0

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F. Cell Diagram:• Cell diagram is the conventional notation for

representing galvanic cells. For example:Zn(s)|ZnSO4(aq,1.00M)|KCl(sat’d)|CuSO4(aq,1.00M)|Cu(s)where:

-the vertical lines represent phase boundaries; -the Zn bar is a solid and the ZnSO4 is a solution, thus a line is drawn in between to show the phase boundary. -the concentration of the species in solution is given, -the salt bridge is included in between two vertical lines.

• The anode is written first at the left (always by convention) and the other components appear in the order in which they are encountered in moving from anode to cathode.

G. The Nernst Equation:

E = E0 - ln QRTnF

(1)Where:

Eºcell is the standard emf of the cell, R is the universal gas constant, T is the absolute temperature,n is the number of moles of electrons transferred during the course of the reaction, F is faraday’s constant, Q is the reaction quotient, given by:

As mentioned earlier, the emf of a cell depends on the concentrations of reactants and products. For a redox reaction of the type a A + b B � c C + d Dthe relationship between emf and concentrations of the species is given by

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• Equation (1) is known as the Nernst equation.

• At 298 K (25ºC), the Nernst equation can be written as

- 0.0257 Vn ln QE

0E =

- 0.0592 Vn log QE0E =

[ ] [ ][ ] [ ]ba

dc

BADC

Q = (2)

(3)

H. Concentration Cells:• Same couple but different concentrations:

Anode M � M2+(L) + 2e-Cathode M2+(R) + 2e- � M

__________________M + M2+(R) � M2+(L) + M

[ ][ ])(

)(ln

20257.0

2

20

RMLM

EE +

+

−=

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I. Spontaneity of Redox Reactions:• When E0 of a given cell is positive, the redox

reaction -taking place in the cell- is spontaneous.

• E0 > 0 ⇔ spontaneous reaction

• To predict whether a reaction would proceed spontaneously as written, one can find its cell potential E and then if it is positive then the reaction is spontaneous; Otherwise, if E is negative then the reaction is not spontaneous as written, the reverse reaction is spontaneous.

Practice Exercises1.Calculate E° and E (the cell emf) for the following cell

reactions:a) Mg(s) + Sn2+(aq) � Mg2+(aq) + Sn(s)

[Mg2+] = 0.045 M; [Sn2+] = 0.035 M.b) 3Zn(s) + 2Cr3+(aq) � 3Zn2+(aq) + 2 Cr(s)

[Cr3+] = 0.010 M; [Zn2+] = 0.0085 M.Answer: (for E)a) + 2.23 V b) = + 0.04 V2. Which species in each pair is a better reducing agent

under standard-state conditions?a) Fe2+ or Ag b) Br- or Co2+

Answer:a) Fe2+ b) Br-

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III. Procedure:A. Galvanic Cells:1. Prepare the following Galvanic cell:

Zn(s)|ZnSO4(0.100 M)||CuSO4(0.100 M)|Cu(s)

- Use 50 mL beakers for the solutions compartments, and Zn and Cu metal bars for the electrodes.

- Obtain from the store room a KCl salt bridge (polyethylene tube containing a 1.00 M KCl in agar gel), and use it for the connection of the two solutions.

- Measure the cell potential. Compare the predicted cell potential from the literature values of the two reduction potentials. Enter data in Table 1.

- Repeat the same experiment by changing the concentrations of the two solutions as follows: a- 0.100 M ZnSO4/0.00100 M CuSO4,b- 0.00100 M ZnSO4/0.100 M CuSO4.

- Compare with the cell potentials calculated from the Nernst equation. Enter values in Table 2.

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2. Prepare the following cell:Zn(s)|ZnSO4(0.100M)||Pb(NO3)2(0.100M)|Pb(s)

-Repeat the procedure in (1) for this cell.

3. Prepare the following cell:Zn(s)|Zn(NO3)2(0.100M)||AgNO3(0.100M)|Ag(s)

-Repeat the procedure in (1) for this cell.

4. Prepare the following cell:Cu(s)|Cu(NO3)2(0.100M)||AgNO3(0.100M)|Ag(s)

-Repeat the procedure in (1) for this cell.

B. Concentration Cells:

- Select a metal M and construct the following concentration cell: M(s)|M2+(0.100M)||M2+(0.0010M)|M(s)

- Measure the cell potential. Compare with the cell potential calculated from the Nernst equation. Enter values in Table 3.