EXP 2 - electrochemical cell PDF

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OBJECTIVES To construct electrochemical cells. To determine the net cell potentials for three electrochemical cells To calculate the net cell potential based on the potentials of the half-reactions that occur and to compare your experimental and calculated values. To measure the Gibbs energy based o...

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OBJECTIVES 1. To construct electrochemical cells. 2. To determine the net cell potentials for three electrochemical cells 3. To calculate the net cell potential based on the potentials of the half-reactions that occur and to compare your experimental and calculated values. 4. To measure the Gibbs energy based on the cell potentials. INTRODUCTION Electrochemistry is a study of a relationship between a chemical reaction and also electricity. It holds the potential to develop a cell that is work on a chemical system by driving an electric current through the system. The cells are identified as electrolytic cells that moved through electrolysis. Electrolysis is used to carry an oxidation-reduction or called a redox reaction which takes place in electrochemical cells. The reaction flows in a direction which it does not occur spontaneously through the movement of electric current in the system while doing work on the chemical system itself and therefore is non-spontaneous. The type of electrochemical cell includes the galvanic cell, electrolytic cell and concentration cell. The galvanic cells and electrolytic cells both are linked in the redox reaction where the anode and cathode electrode are representing for oxidation and reduction process respectively. The electrons for both cells flow from anode to cathode and the salt bridge is used as the passage of ions to maintain electrical neutrality. Theoretically in the galvanic cell, the direction of the electron flow is spontaneous which generates electrical energy which is from chemical energy to electrical energy. So in the cell reaction, the difference in chemical potential energy between higher energy reactants and lower energy products is converted into electrical energy. However, the direction of electron flow in electrolytic cells may be reversed from the direction of spontaneous electron flow in galvanic cells. In other words, the electrolytic cell converted electrical energy into chemical energy. Although the direction of electron flow is reversed, the definition of cathode and anode is still the same where at anode electrode where the oxidation process occurs while the cathode is the place for the reduction process to occur.

APPARATUS 1. Test tubes 2. Pipettes 3. Standard flasks 4. High resistance voltmeter 5. Alligator clips CHEMICALS 1. 0.50M potassium chloride solution 2. 0.10M zinc (II) sulphate solution 3. 0.20M copper (II) sulphate solution 4. 0.50M iron (II) ammonium sulfate solution 5. Zinc metal strip 6. Copper metal strip 7. Iron strip PROCEDURE A salt bridge was prepared for each electrochemical cell by completely dripping a filter paper 0.50M KNO3 solution. Part I: Constructing the zinc-copper electrochemical cell 1. Two test tubes were filled to about ¾ full with 0.10M zinc (II) sulfate solution and 0.20M copper (II) sulphate solution. 2. Strips of zinc and copper were cleaned with steel wool. 3. Alligator clip was used to clip copper strip to one voltmeter terminal and another alligator was used to connect the zinc strip to the voltmeter terminal. The copper strip was dipped into a test tube containing copper solution and zinc strip was dipped into a test tube containing the zinc solution at the same time. The wires were made sure to not come in contact with the

solution. Salt bridge was inserted into both test tubes. The voltage was observed and saltbridge immediately. If the reading on the voltmeter is negative, the connections were reserved to the metal strips and salt-bridge was dipped again. 4. Three readings of voltage were recorded and the anode and cathode of the cell were identified. 5. The temperature at which the measurement was taken was recorded. Part II: Constructing the iron-copper electrochemical cell 1. A cleaned test tube was filled with ¾ full of 0.50M iron (II) ammonium sulfate solution. 2. The iron strips were cleaned with steel wool. The iron strips were connected to one terminal of the voltmeter as before. The copper strip from Part I was kept attached to the other terminal. 3. The metal strips were dipped into the proper solution and connected with the salt-bridge as before. Three positive readings were recorded and anode and cathode of the cell were identifies. 4. The temperature at which the measurement was taken was recorded. Part III: Constructing the iron-zinc electrochemical cell 1. The iron strip attached to the voltmeter stays, copper strip in Part II was replaced with zinc strips in Part I. 2. Metal strips were dipped into proper solution and connected with salt-bridge. Three positive readings were recorded and anode and cathode of the cell were identified. 3. The temperature at which the measurement was taken was recorded.

RESULTS AND CALCULATIONS

Electrolytes

Concentration (M)

Zn2+

0.10

Cu2+

0.20

Fe2+

0.50

Cell No. Cell symbol/notation

Zinc-copper cell Zn(s) | Zn2+(aq) ||

Iron-copper cell

Iron-zinc cell

Fe(s) | Fe2+(aq) || Cu2+ Zn(s) | Zn2+(aq) || Fe2+

Cu2+(aq) | Cu(s)

(aq) | Cu(s)

(aq) | Fe(s)

1. 0.98V

1. 0.62V

1. 0.62V

2. 1.03V

2. 0.63V

2. 0.63V

3. 1.04V

3. 0.67V

3. 0.67V

Average Ecell

1.02V

0.64V

0.35V

T(K)

301.15

301.15

301.15

Ecell

o

∆ G =−nFE ΔG

o

o

∆ G =−nFE

o

o

∆ G =−nFE

o

¿−2(96500 )( 1.02)

¿−2(96500 )(0.64)

¿−2(96500 )(0.35 )

¿−196860 J

¿−123520 J

¿−67550 J

Zinc-Copper Electrochemical Cell

Anode (oxidation)

Cathode (reduction)

Half reaction

−¿ 2+¿ ( aq )+ 2e ¿ Zn ( s ) → Zn¿

−¿ → Cu(s) 2+¿ ( aq )+2e ¿ ¿ Cu

Eo

−0.76 V

+0.34 V

Overall reaction

2+¿ ( aq )+Cu( s) 2+¿ ( aq) → Zn¿ Zn ( s ) +Cu¿

Cell symbol/cell notation

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) o o o Ecell =Ecathode −Eoanode=Eox −Eo¿

¿+0.34−(−0.76 )

Standard cell potential

¿ 1.1V

Half reaction: −¿ →Cu ( s ) 2+¿ ( aq )+ 2e ¿ ¿ Cu

+0.34V

−¿ → Zn ( s ) ¿ 2+¿ ( aq )+ 2e ¿ Zn

-0.76V

Solution:

−¿ ¿ 2+¿ ( aq )+ 2e ¿ Cu

→

Zn ( s )

→

Cu ( s )

+0.34V

2+¿ ( aq ) ¿ ( ) 2+¿ aq + Zn ( s) → Cu ( s )+Zn ¿ Cu

−¿ ¿ 2+¿ ( aq )+ 2e ¿ Zn

-0.76V Ecell =0.34 V −(−0.76 )=1.1V

Equilibrium constant, K: o

Ecell =

RT ln K nF

1.1V =

0.0592V log K 2

log K=

1.1 V × 2 0.0592 V

log K=37.16 K=1.4 × 10

37

Change in free energy, ΔG: ∆ G o =−nFE ocell V ∙ mol e−¿ ×0.93 V e−¿ 96.5 kJ 2 mol × ¿ mol rxn ¿− ¿ 2

¿−1.8 × 10 kJ /mol

Iron-Copper Electrochemical Cell

Anode (oxidation)

Cathode (reduction)

Half reaction

−¿ ¿ 2+¿ ( aq )+ 2e Fe ( s ) → Fe ¿

−¿ → Cu(s) ¿ 2+¿ ( aq )+ 2e ¿ Cu

Eo

−0.44 V

+0.34 V

Overall reaction

2+¿ ( aq )+Cu( s) ¿ 2+¿ ( aq) → Fe ¿ Fe ( s ) +Cu

Cell symbol/cell notation

Fe(s) | Fe2+(aq) || Cu2+(aq) | Cu(s)

Standard cell potential

o

o

o

o

o

Ecell =Ecathode −E anode=Eox −E¿ ¿+0.34−(−0.44 )

¿ 0.78 V

Half reaction: −¿ → Cu(s) ¿ 2+¿ ( aq )+ 2e ¿ Cu

+0.34V

−¿ → Fe ( s ) ¿ 2+¿ ( aq )+ 2e ¿ Fe

-0.44V

Solution:

−¿ ¿ 2+¿ ( aq )+ 2e ¿ Cu

→

Fe ( s )

→

+0.34V

2+¿ ( aq) ¿ ( ) aq 2+¿ +Fe ( s )→ Cu ( s )+Fe ¿ Cu

Equilibrium constant, K: o

Ecell =

RT ln K nF

0.78 V = log K=

0.0592V log K 2

0.78 V ×2 0.0592 V

Cu ( s )

−¿ ¿ 2+¿ ( aq )+ 2e ¿ Fe

-0.44V Ecell =0.34 V −(−0.44 )=0.78 V

log K=26.35 K=2.2 ×10

26

Change in free energy, ΔG: ∆ G o =−nFE ocell V ∙ mol e−¿ ×0.78 V e−¿ 96.5 kJ 2 mol × ¿ mol rxn ¿− ¿ 2

¿−1.5 × 10 kJ /mol

Iron-Zinc Electrochemical Cell

Anode (oxidation)

Cathode (reduction)

Half reaction

−¿ 2+¿ ( aq )+ 2e ¿ Zn ( s ) → Zn¿

−¿ → Fe(s) 2+¿ ( aq )+2e ¿ ¿ Fe

Eo

−0.76 V

−0.44 V

Overall reaction

2+¿ ( aq )+Fe (s ) 2+¿ ( aq) → Zn¿ Zn ( s ) + Fe¿

Cell symbol/cell notation

Zn(s) | Zn2+(aq) || Fe2+(aq) | Fe(s)

Standard cell potential

o

o

o

o

o

Ecell =Ecathode −E anode=Eox −E¿

¿−0.44−(−0.76 ) ¿ 0.32V

Half reaction: −¿ → Fe(s) ¿ 2+¿ ( aq )+ 2e ¿ Fe

- 0.44 V

−¿ → Zn ( s ) 2+¿ ( aq )+ 2e ¿ ¿ Zn

-0.76V

Solution:

−¿ 2+¿ ( aq )+ 2e ¿ ¿ Fe

→

Zn ( s )

→

-0.44V

2+¿ (aq ) ( ) aq 2+¿ +Zn ( s) → Fe( s )+Zn¿ ¿ Fe

Equilibrium constant, K: o

Ecell =

RT ln K nF

0.32V = log K=

0.0592V log K 2

0.32 V × 2 0.0592 V

log K=10.81

Fe ( s )

−¿ ¿ 2+¿ ( aq )+ 2e ¿ Zn

-0.76V Ecell =−0.44 V −(−0.76 )=0.32V

10

K=6.5 × 10

Change in free energy, ΔG: o

o

∆ G =−nFEcell

V ∙ mol e−¿ ×0.32 V e−¿ 96.5 kJ 2 mol × ¿ mol rxn ¿− ¿ ¿−6.2 × 101 kJ /mol

DISCUSSIONS A galvanic cell was set up by connecting a two half-cell using a salt bridge. The salt bridge used to connect between the two half-cell using a filter paper soaked in 0.1M KMnO 4 and this will allow the electrons flows between the electrolytes. The reduction process was occurred at cathode while oxidation occur at anode. The oxidation means that a process which the reactant lose its electron while reduction process gain its electron. Electrons will flow from anode which is a negative terminal to cathode which is a positive terminal resulting electrical potential detected using voltmeter.

In zinc-copper electrochemical cell, a half-cell of containing 0.20M copper (II) sulphate solution connected with a half-cell containing 0.10M zinc (II) sulphate solution using a salt bridge. The cathode in this reaction is copper strip dipped into the copper (II) sulphate solution where reduction process occurs while the anode is zinc strips dipped into zinc (II) sulphate

solution where the oxidation occurs. Standard electrode potential of zinc which is -0.76V is more negative compared to copper which is +0.34V and resulting in zinc act as anode. The theoretical electrical potential between two electrolytes is 1.1V while the electrical potential collected from the experiment is 1.0V makes the experimental result is lower from the theoretical value.

In iron-copper electrochemical cell, a half-cell containing 0.50M iron (II) ammonium sulphate solution connected with a half-cell containing 0.20M copper (II) sulphate solution using a salt bridge. The cathode in this reaction is copper strip dipped into copper (II) sulphate solution where reduction process occur while the anode is iron strip dipped into iron (II) ammonium sulphate solution where the oxidation occur. Standard electrode potential of iron is -0.44V which is more negative compared to copper which is +0.34V and resulting in iron strip act as anode. The theoretical electrical potential between the two electrolytes is 0.78V while the electrical potential collected from the experiment is 0.64V. And the outcome it the experimental value is lower compared to the theoretical value.

In iron-zinc electrochemical cell, a half-cell containing 0.10M zinc (II) sulphate solution connected with a half-cell containing 0.50M iron (II) ammonium sulphate solution using a salt bridge. The cathode in this reaction is iron strip dipped into iron (II) ammonium sulphate solution where reduction process occur while the anode is zinc strip dipped into zinc (II) sulphate solution where the oxidation occur. Standard electrode potential of zinc is -0.76V which is more negative compared to zinc which is -0.44V and resulting in zinc strip act as anode. The theoretical electrical potential between the two electrolytes is 0.32V while the electrical potential collected from the experiment is 0.35V.

The possible sources of error that can affect the result of this experiment is maybe because by connection of circuit that was connected properly during the experiment. So precaution step must be followed to ensure the connection is connected properly. Another possible error is when polishing the electrodes. Electrodes must be polished to avoid any impurities during the experiment. Lastly, the instability of voltmeter scale was probable cause of the error in the value.

CONCLUSION

As a conclusion, net cell potential for three electrochemical cells which are zinc-copper electrochemical cell, iron-copper electrochemical cell and iron-zinc electrochemical cell were determined. Experimental value of Zinc-copper electrochemical cell and iron-copper electrochemical cell were lower compared to the theoretical value due to several reasons. From calculated net cell potential, Gibbs energy can be determined. REFERENCES Angel

C.

de

Dios.

(2010).

Introduction

to

Electrochemistry.

Retrieved

from:

https://bouman.chem.georgetown.edu/S02/lect25/lect25.htm

Bockris O. J., & Despić R. A. (2011). Electrochemical Reaction. Retrieved from: https://www.britannica.com/science/electrochemical-reaction

The Student Room website. (2007). Effect Of Concentration On Electrode Potential. Retrieved from https://www.thestudentroom.co.uk/showthread.php?t=360208

QUESTIONS 1. What is the reason for using a salt bridge? The salt bridge completes the electrical circuit and allows ions to flow through both halfcells. 2. Calculate the standard cell potential of a cell constructed from Mg 2+/Mg and Ni2+/Ni. Which is the anode and which is the cathode?

−¿ → Mg(s) ¿ 2+¿ ( aq )+ 2e ¿ Mg

Eo = -2.37V

−¿ →∋( s ) ¿ 2+¿ ( aq )+ 2e ¿ ¿

Eo = -0.25V

o o o o o Ecell =Ecathode−E anode=Eox−E¿

¿−0.25−(−2.37 ) ¿+2.12 V

2+¿ /Mg ∴ Anode= Mg¿ 2+¿ /¿ ∴Cathode=¿ ¿

3. Using the Nernst Equation, what would be the potential of a cell with [Ni 2+] = [Mg2+] = 0.10M? Half reaction: −¿ → Mg(s) ¿ 2+¿ ( aq )+ 2e ¿ Mg −¿ →∋( s ) ¿ 2+¿ ( aq )+ 2e ¿ ¿

Solution:

Eo = -2.37V

Eo = -0.25V

−¿ 2+¿ ( aq ) + 2e ¿ ¿¿

¿ ( s)

Mg ( s )

−¿ ¿ 2+¿ ( aq )+ 2e ¿ Mg

2+¿ ( aq) ¿ ( ) 2+¿ aq + Mg ( s )→∋ ( s ) +Mg ¿¿

-0.25V

-2.37V Ecell =−0.25 V −(−2.37 V )=2.12V

2 +¿

¿ Mg2+¿ ¿ Q =¿ ¿

0.10 M 0.10 M

¿1

o E=E cell −

¿ 2.12−

RT lnQ nF

(8.31 ×298) ln 1 2(96500)

¿0

Change in free energy, ΔG: ∆ G o =−nFE ocell V ∙ mol e−¿ ×2.12 V e−¿ 96.5 kJ 2 mol × ¿ mol rxn ¿− ¿ 2

¿−4.1 ×10 kJ /mol

4. What is the information deduced from the sign and the magnitude of ΔE and ΔG from this experiment? ΔE > 0

ΔG < 0 It is a non-spontaneous reaction. Redox reaction is negative....