Galvanic Cells, the Nernst Equation (Experiment 4) PDF

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Galvanic Cells, the Nernst Equation (Experiment 4)ABSTRACTGalvanic cells or also known as voltaic cells is one of the types of electrochemical cells which is a spon taneous reaction ofoxidation-reduction reaction occurs that will produce electrical energy. The standard cell potential (E°cell) is the...

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Galvanic Cells, the Nernst Equation (Experiment 4) ABSTRACT Galvanic cells or also known as voltaic cells is one of the types of electrochemical cells which is a spontaneous reaction of oxidation-reduction reaction occurs that will produce electrical energy. The standard cell potential (E°cell) is the measurement of the potential difference of the two (2) half cells, specifically the cathode and the anode. The objectives of this experiment is to first determine the reduction potentials of given redox couples, second, to understand the movement of electrons, anions and cations in a galvanic cell and to study the factors that affects the cell potential of a galvanic cell and to estimate the concentration of ions in solution by the Nernst equation. With prepared solutions for the two parts of the experiment, the cell potentials were measured 0.737 V, 1.042 V, and 0.231 V respectively. The measured values were gathered and recorded. The group was able to analyze the effects of the difference in the Molarity to the cell potential value of a reaction. It was found out that the difference in the Molarity makes the reaction happen. The objectives of the experiment were achieved. The relative reduction potentials for the redox pairs given were measured and the concentration of ions in the four (4) unknown solutions were estimated by using the Nernst equation. Keywords: Nernst equation, anode, cathode, cell potential, oxidation-reduction

INTRODUCTION Galvanic cells or also known as voltaic cells is one of the types of electrochemical cells which is a spontaneous reaction of oxidation-reduction reaction occurs that will produce electrical energy. Oxidation-reaction is a chemical reaction wherein electrons from one molecule or atom are transferred to another. It occurs by separating the redox reactions into two (2) half-cell reactions namely, the anode and the cathode. The anode is the site for oxidation where the substance loses its electrons. The anode is connected to the cathode, the site for reduction. (Flowers et al, n.d) The standard cell potential (E°cell) is the measurement of the potential difference of the two (2) half cells, specifically the cathode and the anode. The cell potential allows the flow of electron from one cell to another. The standard cell potential (E°cell) is computed by subtracting the cell potential of anode from the cell potential of cathode which are both in volts unit, under standard conditions (Libretext, 2016). The Nernst equation is useful in determining the cell potential under nonstandard conditions. It relates to the measured standard potential and quotient of reaction resulting to accurate constants (Libretext, 2017). The objectives of this experiment is to first determine the reduction potentials of given redox couples, second, to understand the movement of electrons, anions and cations in a galvanic cell and to study the factors that affects the cell

potential of a galvanic cell and to estimate the concentration of ions in solution by the Nernst equation. MATERIALS AND METHODS Three (3) pairs of known chemicals were given as the main objects in the experiment. The chemicals were prepared prior to the time assigned for the experiment. There were two (2) set-ups fir the experiment. The groups were given five (5) minutes for the first set up to measure the value of the cell potential of each of the galvanic cells connected by a salt bridge. Another five (5) minutes was given for the measuring of the redox couple. A multi-tester was used in this method. With the alligator clip was connected to the Copper and Zinc strips. The opposite side of the alligator clips were connected to the multi-meter for the reading of the cell potentials of the given chemicals. The data needed for computations were then recorded.

Table 3. Reduction Potencials of Unknown Redox Pairs Galvanic Cell

Ecell measured

Cu-Zn

0.737 V

Reduction Potential (Experimental) -0.053

Reduction Potential (Theoretical) 0.31 V

Zn-Fe

0.231 V

-0.559

-0.47

Zn-Zn

0V

-0.79

-0.79

% Erro r 117. 1% 18.9 % 0%

Table 3 contains the data gathered from the first part of the experiment. From the value of Ecell, the experimental value of the reduction potential for the galvanic cell Cu-Zn and Zn-Fe. The computed percent error is also in the table. The computations for the values in the experiment are shown in the calculations part of the report.

Figure 1. Experimental Set Up 1

Table 4. The Nernst Equation of an Unknown Concentration

Figure 2. Experimental Set Up 2 RESULTS AND DISCUSSION The measured values from the experiment are shown in table 1. The equations for the anode half reaction and cathode half reaction are also shown in table 2. These are derived from the given redox pairs. Table 1. Reduction Potentials of the Redox Pairs Galvanic Cell Cu-Zn Zn (Anode) ; Cu (Cathode) Cu-Fe Fe (Anode) ; Cu (Cathode) Zn-Fe Zn (Anode); Fe (Cathode)

Ecell measured 0.737 V 1.042 V 0.231 V

Table 2. Equations for Half-Reaction Anode Half-Reaction 2+ + 2𝑒 − 𝑍𝑛(𝑠) → 𝑍𝑛(𝑎𝑞) 2+ + 2𝑒 − 𝐹𝑒(𝑠) → 𝐹𝑒(𝑎𝑞) 2+ + 2𝑒 − 𝑍𝑛(𝑠) → 𝑍𝑛(𝑎𝑞)

Cathode Half-Reaction 2+ + 2𝑒 − → 𝐶𝑢(𝑠) 𝐶𝑢(𝑎𝑞) 2+ + 2𝑒 − → 𝐶𝑢(𝑠) 𝐶𝑢(𝑎𝑞) 2+ + 2𝑒 − → 𝐹𝑒(𝑠) 𝐹𝑒𝑎𝑞)

Solutio n

Concentratio n of Cu(NO3)

1 2 3 4

0.1 mol/L 0.001 mol/L 0.00001 mol/L 0.0000001 mol/L

Ecell Experimenta l 21.9 V 25.9 V 27.9 V 95.2 V

-log, [Cu2+] , pCu 1 3 4 7

Ecell Calculate d 21.87 V 25.81 V 27.75 V 94.99 V

Table 4 shows the data gathered from the second part of the experiment. The value of the Ecell is the measured value from the given unknown concentration. From the Nernst equation, the values for -log, [Cu2+], pCu and the calculated value of Ecell were determined for the 4 unknown solutions. The computations for the values in the experiment are shown in the calculations part of the report. The formulas used in computing for the values needed in the experiment are shown in the following equations:

𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑐𝑎𝑡ℎ𝑜𝑑𝑒 − 𝐸𝑎𝑛𝑜𝑑𝑒 0.0592 ° − 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑐𝑒𝑙𝑙 log 𝑄 𝑛 𝑀 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠

𝑄 = 𝑀 𝑜𝑓 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠

Where: Ecell – standard cell potential Ecathode – cell potential of cathode Eanode – cell potential of anode n – no. of electrons exchanged Q – reduction quotient M – molarity

Eqn. 1 Eqn. 2 Eqn 3.

CONCLUSION AND RECOMMENDATIONS The objectives of the experiment were achieved. The relative reduction potentials for the redox pairs given were measured and the concentration of ions in the four (4) unknown solutions were estimated by using the Nernst equation. The concentration affects the cell potential of a cell especially on the electrode where the concentration was altered. The Nernst equation is useful in determining the cell potential under non-standard reaction with different molarities. Galvanic cell was also used to measure the solubility of product of a soluble substance and to compute for the concentration. There were errors encountered while performing the experiment such as the accuracy of the data gathered from the different redox pairs and unknown solutions since the containers (beakers) were too small and the electrodes are touching the salt bridge. The electrodes are not properly immersed in the solutions. Another error is the accuracy of the data given by the multi-meter since there were no stable or constant value displayed when the cell potential was measured. It is recommended to use a bigger container for the solutions to avoid the touching of alligator clips and electrodes with the salt bridge. It is also recommended to have a precise/accurate data by waiting for the multi-meter to flash a constant value of the cell potential.

REFERENCES Flowers, P. et al., (n.d.). Chemistry. Retrieved from https://opentextbc.ca/chemistry/chapter/17-2-galvanic-cells/ Libretexts, (2016). The Cell Potential. Retrieved from https://chem.libretexts.org/Textbook_Maps/Analytical_Che mistry/Supplemental_Modules_(Analytical_Chemistry)/Elect rochemistry/Voltaic_Cells/The_Cell_Potential Libretexts. (2017). Nernst Equation. Retrieved from https://chem.libretexts.org/Textbook_Maps/Analytical_Che mistry/Supplemental_Modules_(Analytical_Chemistry)/Elect rochemistry/Nernst_Equation

B. CALCULATIONS ( from table 4 ) APPENDIX Concentration of Cu(NO3)2 A. CALCULATIONS (from Table 3) Solution Number 2: 𝐶1 𝑉1 = 𝐶2 𝑉2

1. Reduction Potencial (experimental) for Cu-Zn Cu-Zn:𝐸𝐶𝑢2+ /𝐶𝑢 = 𝐸𝑐𝑒𝑙𝑙 𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑑 + (−0.79 𝑉) Cu-Zn:𝐸𝐶𝑢2+ /𝐶𝑢 = 0.737 𝑉 + (−0.79 𝑉)

(0.1

𝑚𝑜𝑙 𝐿

)(1 𝑚𝑙) = 𝐶2 (100 𝑚𝑙) 𝑪𝟐 = 𝟎. 𝟎𝟎𝟏 M

Cu-Zn:𝐸𝐶𝑢2+ /𝐶𝑢 = −0.053 𝑉

Solution Number 3:

2. Reduction Potencial (experimental) for Zn-Fe Zn-Fe:𝐸𝐹𝑒2+ /𝐹𝑒 = 𝐸𝑐𝑒𝑙𝑙 𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑑 + (−0.79 𝑉) Zn-Fe:𝐸𝐹𝑒2+ /𝐹𝑒 = 0.231 𝑉 + (−0.79 𝑉)

(0.001 M)(1 ml) = 𝐶3 (100 𝑚𝑙 ) 𝑪𝟑 = 𝟎. 𝟎𝟎𝟎𝟎𝟏 M Solution Number 4: (0.00001 𝑀)(1 𝑚𝑙) = 𝐶4 (100 𝑚𝑙 )

Zn-Fe:𝑬𝑭𝒆𝟐+ /𝑭𝒆 = −𝟎. 𝟓𝟓𝟗 𝑽

𝑪𝟒 = 𝟎. 𝟎𝟎𝟎𝟎𝟎𝟎𝟏 𝑴 3. Percent Error for Cu-Zn % 𝑒𝑟𝑟𝑜𝑟 = |

𝑉𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 −𝑉𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙

% 𝑒𝑟𝑟𝑜𝑟 = |

−0.053 𝑉−0.31 𝑉 |x100 0.31 𝑉

𝑉𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙

|x100

-log,[𝐶𝑢 2+]= -log (0.001)

% 𝒆𝒓𝒓𝒐𝒓 = 𝟏𝟏𝟕. 𝟏 %

-log,[𝑪𝒖𝟐+]= 3

4. Percent Error for Zn-Fe % 𝑒𝑟𝑟𝑜𝑟 = |

𝑉𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 −𝑉𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙

% 𝑒𝑟𝑟𝑜𝑟 = |

−0.559 𝑉−(−0.47 𝑉) |x100 −0.47 𝑉

𝑉𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙

% 𝒆𝒓𝒓𝒐𝒓 = 𝟏𝟖. 𝟗 %

-log,[𝐶𝑢 2+], 𝑝𝐶𝑢 for solution number 2:

-log,[𝐶𝑢 2+], 𝑝𝐶𝑢 for solution number 3: |x100

-log,[𝐶𝑢 2+]= -log (0.00001) -log,[𝑪𝒖𝟐+]= 5 -log,[𝐶𝑢 2+], 𝑝𝐶𝑢 for solution number 4: -log,[𝐶𝑢 2+]= -log (0.0000001) -log,[𝑪𝒖𝟐+]= 7

𝐸𝑐𝑒𝑙𝑙 𝑐𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑 for solution number 1: 𝐸° = 𝐸𝑐𝑒𝑙𝑙 −

0.0592 (𝑝𝐶𝑢) 2

𝐸° = 21.9 𝑉 −

0.0592 (1) 2

𝑬° = 𝟐𝟏. 𝟖𝟕 𝑽 𝐸𝑐𝑒𝑙𝑙 𝑐𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑 for solution number 2: 𝐸° = 𝐸𝑐𝑒𝑙𝑙 −

0.0592 (𝑝𝐶𝑢) 2

𝐸° = 25.9 𝑉 −

0.0592 (3 ) 2

𝑬° = 𝟐𝟓. 𝟖𝟏 𝑽 𝐸𝑐𝑒𝑙𝑙 𝑐𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑 for solution number 3: 𝐸° = 𝐸𝑐𝑒𝑙𝑙 −

0.0592 (𝑝𝐶𝑢) 2

𝐸° = 27.9 𝑉 −

0.0592 (5) 2

𝑬° = 𝟐𝟕. 𝟕𝟓 𝑽 𝐸𝑐𝑒𝑙𝑙 𝑐𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑 for solution number 4: 𝐸° = 𝐸𝑐𝑒𝑙𝑙 −

0.0592 (𝑝𝐶𝑢) 2

𝐸° = 95.2 −

0.0592 (7) 2...