Galvanic-Cells - Galvanic Cells Lab Report PDF

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Galvanic Cells Lab Report...

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Mapúa University Intramuros, Manila S.Y. 2019 - 2020

Experiment No. 1:

Galvanic Cells, the Nernst Equation

Name: Sumande, Cedrix V.

CE -1

Section: CM011L

Group no. 3:

Date Performed:

Sumande, Cedrix

January 16, 2020 Escala, Jc

Simons, Mark

Date of Submission: January 23, 2020

Dela Cruz, Kenneth Teodosio, Tsidkenu Rophi M.

Elizabeth Espiritu Instructor

GRADE

Experiment No. 1:

Galvanic Cells, the Nernst Equation INTRODUCTION: Electrochemical cells are two types, galvanic and electrolytic, both employing the principle of oxidation-reduction (redox) reaction. In galvanic (or voltaic) cells (this experiment), redox reactions occur spontaneously as is common with all portable batteries of which we are very familiar. Electric cars, flashlight, watches, and power tools operate because of a specific spontaneous redox reaction. Electrolytic cells (experiment 33) are driven by non-spontaneous redox reactions, reactions that require energy occur. The recharging of batteries, electroplating and refining of metals, and generation of various gases all require the use of energy to cause the redox reaction to proceed. Experimentally, when copper wire is placed into a silver ion solution (see opening photo), copper atoms spontaneously donate electrons (copper atoms are oxidized) to the silver ions (which are reduced). Silver ions migrate to the copper atoms pick up electrons and form silver atoms at the copper metal-solution interface; the copper ions that form then move into the solution away from the interface. The overall reaction that occurs at the interface is: Cu(s) + 2 Ag+(aq) forming 2Ag (s) + Cu2+ (aq) This redox reaction can be divided into an oxidation and a reduction half-reaction. Each half-reaction, called a redox couple, consist of the reduced state and the oxidized state of the substance.

OBJECTIVE: 1. To measure the relative reduction potentials for a number of redox couples 2. To develop an understanding of the movement of electrons, anions, and cations in a galvanic cell 3. To study factors affecting cell potentials 4. To estimate the concentration of ions in solution using the Nernst equation

DATA & RESULT A. Reduction Potentials of Several Redox Couples Galvanic cell

Ecell measured

Anode

Equation for Cathode Anode HalfReaction

Cu-Zn

0.899 V

Zn

Zn 2e-

Zn2+ + Cu

Equation for Cathode Half-Reaction Cu2+ + 2eCu

Cu-Mg

-----------

_______

___________

_________

_________

Cu-Fe

0.624 V

Fe

Fe Fe2+ + 2e-

Cu

Cu2+ + 2eCu

Zn-Mg

________

_______

___________

_________

_________

Fe-Mg

_______

_______

___________

_________

________

Zn-Fe

0.212 V

Zn

Zn 2e-

1. Write balanced equation for the six cell reactions   

Cu2+ + Zn Cu2+ + Fe Fe2+ + Zn

DATA & RESULT

Cu + Zn2+ Cu + Fe2+ Fe + Zn2+

Zn2+ + Fe

Fe2+ + 2eFe

B. Complete the table Galvanic Cell

Ecell Measured

Redox Couple

Reduction Potential

Reduction Potential

(Experimental )

(Theoretical )

% Error

Cu-Zn

0.899 V

Cu2+/Cu

0.899 V

0.31

0.589 %

Zn-Fe

0.212 V

Fe2+/Fe

0.212 V

-0.47

0.682 %

Zn-Zn

0

Zn2+/Zn

-0.79 V

-0.79

0.00

Zn-Mg

_______

Mg2+/Mg

__________

-2.40

________

ZnUnknown

________

X2+/X

__________

________

________

B. Effect of Concentration Changes on Cell Potential 1. Cell Potential of Concentration cell: 048 mV Anode half-reaction: Cu

Cu2+ (0.01M)+ 2e-

Cathode half-reaction: Cu2+ (1M)+ 2e-

LABORATORY QUESTIONS:

Cu

1. Part A.3. The filter paper salt bridge is not wetted with the 0.1 M KNO3 solution. As a result, will the measured potential of the cell be too high, too low, or unaffected? Explain. 

A difference in charge is formed when electrons exit one half of a galvanic cell and migrate to another. If the filter paper salt bridge is not wetted, the difference in charge will prevent more electrons from flowing. A salt bridge enables ion flow to maintain a balance in charge between the oxidation and reduction vessels while retaining separate contents of each. Electrons can flow again with the charging gap balanced, and the reduction and oxidation reactions can continue. In general, keeping the two cells separate is preferable from the point of view of eliminating variables from an experiment. When no direct contact between electrolytes is allowed, there is no need to make allowance for possible interactions between ionic species.

2. Part A.3. A positive potential is recorded when the copper electrode is the positive electrode. Is the copper electrode the cathode or the anode of the cell? Explain. 

The positive electrode is the cathode, where reduction occurs, and electrons are gained. It is indeed the cathode.

3. Part A.5. The measured reduction potentials are not equal to the calculated reduction potentials. Give two reasons why this might observe 

The measured reduction potential is equal to the estimated reduction potential, it can be due to measuring equipment inaccuracy or due to over-potential and surface activity at the electrodes or ion activities in the solution.

4. Part B.2. Would the cell potential be higher or lower if the NH3(aq) had been added to the 1 M CuSO4 solution instead of the 0.001 M CuSO4 solution of the cell? Explain 

The cell potential will decrease if 1 M CuSO4 is applied instead of 0.001 M CuSO4. This is because there will be a higher concentration of CuSO4 serving as a reducer. This will result in an additional reduction agent that reduces the chances of redox relative to the lower CuSO4 concentration.

5.Part B.3. The cell potential increased (compared to Part B.2) with the addition of the NasS solution to the 0.001 M CuSO4 solution. Explain. 

The cell potential increased with the addition of Na2S solution because more electrons travel from anode to cathode when the 0.001 M CuSO4 solution discharges electrons

6. Part C. As the concentration of the copper (II) ion increased from solution 4 to solution 1, did the measured cell potential increase or decrease? Explain why the change occurred. 

The measured potentials increased as the concentration of copper (II) ion increased since, electrons released from the 0.001 M CuSO4 solution flow from anode to cathode

7. Part C. Suppose the 0.1 M Zn2+ solution had been diluted (instead of the Cu2+ solution), would the measured cell potentials have increased or decreased? Explain why the change occurred 

the spontaneous reaction increases [Zn2+] and decreases [Cu2+]. By diluting the Zn half-cell, it moves the reaction away from equilibrium, which will increase the cell potential, and if by diluting the Cu half-cell the reaction will move towards equilibrium and the cell potential will decreases.

8. Part C. How would you increased or decreased the Cu2+ concentrations and/or increase or decrease the Zn2+ concentration to maximize the cell potential? Explain how the change for each ion would maximize the cell potential. 

Any change to each ion to the system that pushes it further away from equilibrium, causes the potential to go up. Any change to the system that pushes the system towards equilibrium, causes the potential to go down. Therefore, any changes that are made to both Zn2+and Cu2+, the system either moves closer or away from the equilibrium. The reaction quotient for example Q=[Zn2+]/[Cu2+] of equation Zn(s)+ Cu2+(aq)→ Zn2+ (aq)+ Cu(s)of this redox reaction helps to understand the concept. For example, when all concentrations are 1 Mthe Q will equal 1. But this reaction is spontaneous which K > 1. As the cell reaction proceeds and Q approaches K the cell potential drops. Therefore, when both the concentrations of Zn2+and Cu2+are changed Q=[Zn2+]/[Cu2+]) will be affected on their ratio, which will effect on the cell potential also. For example, the potential of the cell increases when the concentration of Cu2+ increases Zn2+is constant at 1.0 M.

CONCLUSION: In this lab, we built various galvanic cells and learned how to measure the cell potential. The Nernst equation tells us the effect of the concentration on cell potential. If we increase the product concentration, the cell potential will decrease and if we decrease the product concentration, the cell potential will increase. In the first part of this experiment, we learned that the experimental cell potentials are not the same as the theoretical cell potentials but are very close. For part three, we learned how the temperature affects cell potentials. Higher temperature increased the cell potential for the Fe-Cu and decreased the cell potential for Zn-Cu. There were a lot of errors that might have occurred such as incorrectly setting up the voltmeter or incorrectly calculating the theoretical cell potential. Which may be why the theoretical values are different from the experimental values....