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EXPERIMENT # 2.2 Galvanic Cells, The Nernst Equation

GENERAL CHEMISTRY FOR ENGINEERS LABORATORY CHM011-L/OL57

ESCOTE, SHANE ANGELO T. 2020117323

Mapúa University October 2020

INTRODUCTION

Electrochemistry is the study of chemical reactions that includes the movement of electrons between solutions. It is also the study of the relationship of chemical reactions and electricity. There are two types of electrochemical cells, galvanic and electrolytic, both employs the principle of oxidation-reduction (redox) reaction. In an oxidation process, the electrons are lost while in the reduction process, electrons are gained. To have a beneficial electrical work, galvanic cells binds electrical energy that are available from electron transportation in a redox reaction. In the experiment, the usage of a galvanic cell was utilized to understand the following objective: (1) Measure relative reduction potentials for several redox couples, (2) Develop depth understanding in the movement of electrons, anions, and cations in a galvanic cell, (3) Study the factors that affects ell potentials, (4) Estimate the concentration of an unknown solutions using the Nernst equation. In a galvanic cell, a redox reaction occurs spontaneously like all portable batteries. This reaction proceeds by splitting the oxidation and reduction solutions whereas the transferring of electrons takes place in an external route called salt bridge. Salt bridge are junctions that connects the anode and cathodes compartments in a cell or an electrolyte solution (Byju’s, n.d.). It is usually made of salt solutions. For example, in the experiment, the chemical reaction of copper (cathode) and iron (anode). The cathode is the location where the cations transfer, and the positive ions go towards positive terminal in order to neutralize electrons gained in the reduction. This can be expressed in half-cell equation of:

EQUATION 1: Fe (s) -> Fe2+ (aq) + 2e- Oxidation half-reaction Cu2+ (aq) + 2e- -> Cu (s) Reduction half-reaction Fe (s) +Cu2+ (aq)

->

Cu (s) + Fe2+ (aq)

The difference in reduction potential of two half-cell of an electrochemical cell are called cell potential (Lumen, n.d.). In every galvanic cell, the cell potentials are measured in volts by using voltmeter. In any case it is impossible to directly measure the potential reduction of the two half-cell in a electrochemical cell however due to constant research of chemists they have formulated a process to measure the ability of reduction in a table of standard reduction potentials (Sparks Notes Editors, n.d.). In the experiment, the reaction of the copper and iron solutions have the cell potential of: EQUATION 2:

E cell

=E Cu 2+

, Cu

–E Fe 2+

, Fe

E cell

=E Cu 2+

, Cu

–E Fe 2+

, Fe

E cell

=E Cu 2+

, Cu

–E Fe 2+

, Fe Ecell = ECu2+, Cu – EFe2+, Fe

Since the copper ion has a greater reduction potential than iron ion does and the cell potential must be a positive value, the copper ion must be placed before the iron ion as expressed in equation 2. The measured cell potential corresponds to the standard cell potential when the concentrations of all ions are 1 mol/L and the temperature of the solutions is 25C°. Otherwise, the Nernst equation can be applied. This equation is expressed as: Since copper has a greater reduction potential than iron the cell potential value must be positive. Hence, the copper (cathode) is placed in front of the iron (anode) as expressed in

equation 2. The cell potential calculated is equal to the standard cell potential when all concentrations are 1 mol/L and with a 25C° temperature. Otherwise, the use of Nernst equation is necessary. The equation is expressed in: EQUATION 3:

Ecell =E °cell−

0.0592 logQ n

Whereas the E°cell is the actual cell potential, n is the number of electrons transferred, Q is the molar concentrations of the product divided by the reactant and raised to the power of its coefficient in the balanced equation. In the whole experiment, three (3) galvanic cells are set up with the apparatus for voltaic cell. The researcher only measured the voltage of each reaction to calculate the unknown concentration of the redox solution. It is composed of copper and zinc ions to produce a redox reaction. Meanwhile, the researcher immersed himself in a virtual laboratory entitled Redox Reactions: Discover how batteries work. Through the laboratory the researcher learned about the essence of periodic table in identifying the oxidizing value of each compound and element, and the way batteries are made to produce electricity. Moreover, the virtual laboratory gave necessary information to carry out the experiment.

RESULTS AND DISCUSSION Table 1.1. Reduction Potentials of Several Redox Couples Anod

Cu-Zn (Ecell = 0.79 V) Equation for Anode Half-Reaction Cathode

Equation for Cathode Half-

e _Zn_

Anod e _Fe_

Anod e _Zn_ _ Galvani c Cell

_Zn(s)___ Zn2+ (aq) + 2e-_

_Cu_

Cu-Fe (Ecell = 0.97 V) Equation for Anode Half-Reaction Cathode _Fe(s)___Fe2+(aq) +_2e-_

_Cu_

Zn-Fe (Ecell = 0.20 V) Equation for Anode Half-Reaction Cathode _Zn(s)___ Zn2+ (aq) + 2e-_

Measured Anod Ecell e

Cu-Zn

0.79 V

Zn

Cu-Fe

0.97 V

Fe

Zn-Fe

0.20 V

Zn

_Fe__

Equation of anode Zn(s) Fe(s) Zn(s)

Reaction _Cu2+(aq) + 2e-___Cu(s)_

Equation for Cathode HalfReaction _Cu2+ (aq) + 2e-___Cu(s)_

Equation for Cathode HalfReaction _ Fe2+(aq) +_2e-___ Fe(s)_

Cathode

Equation of Cathode

Zn2+(aq) + 2e-

Cu

Cu2+(aq) + 2e-

Fe2+(aq) + 2e-

Cu

Cu2+(aq) + 2e-

Cu(s)

Zn2+(aq) + 2e-

Fe

Fe2+(aq) + 2e-

Fe(s)

Oxidation: Zn -> Zn2+ + 2e-

E° = - 0.76V

Reduction: Cu2+ + 2e- -> Cu

E° = 0.34V

Overall: Zn + Cu2+ -> Zn2++ Cu

E° cell= 1.1V

Percent Error: 28.18%

Oxidation: Fe -> Fe2++ 2e-

E° = - 0.44V

Reduction: Cu2+ + 2e- -> Cu

E° = 0. 34V

Overall: Fe + Cu2+ -> Fe2+ + Cu

E° cell = 0.78V

Percent Error: 24.36%

Oxidation: Zn -> Zn2+ + 2e-

E° = -0.79V

Reduction: Fe2+ + 2e- -> Fe

E° = -0.44V

Overall: Zn + Fe2+ -> Zn2+ + Fe

E° cell = 0.35V

Cu(s)

Percent Error: 42.86%

Using the equation for cell potential, in which the highest standard reduction potential (cathode) is subtracted by the lowest standard reduction potential (anode) to calculate the voltage release by the redox couple, relate to equation 2. The researcher analyzed that the measured and actual value of the Ecell produces a high percentage error. This may be caused by the impurities of the solution during the experiment. Furthermore, it maybe because of a systematic error using the laboratory equipment and an error in the amount of concentration used. Table 1.2. Effect of Concentration Changes on Cell Potential Cu | Cu2+(aq,0.001M) || Cu2+(aq,1M) | Cu Cell Potential of Concentration Cell: 0.087 V Equation for Anode Half-Reaction Equation for Cathode Half-Reaction _ Cu(s) Cu2+(aq, 0.001M) + 2e- _ _ Cu2+(aq, 1M) + 2e-___Cu(s)_ In the calculation of the redox couple, it is concluded that the ionization of the anode substance to transfer its electron to the cathode produces a spontaneous cell potential of 0.087V. According to Khan Academy, the low cell potential value is because of the difference in the concentration of the substances. It includes the principle of concentration cells. A concentration cell is an electrolytic cell that consist of two half-cells with similar electrodes, but different concentration. It acts to dilute the higher concentrated solution and concentrate the more dilute solution to create a voltage when the cell reaches an equilibrium. Table 1.3. The Nernst Equation and an Unknown Concentration Concentration Cell Solution 1 – Solution 2 Solution 1 – Solution 3 Solution 1 – Solution 4

Ecell measured (V) 0.06 0.12 0.18

Concentration of the Unknown, M 1.38 x 1034 M 1.29 x 1032 M 1.20 x 1030 M

Constants: Temperature = 25oC [Solution 1] = 0.1 M Cu2+ (consider Solution 1 as the concentrated solution and the cathode) Since the principle of the cell potential was not met, the used of Nernst equation was necessary. The Nernst equation enables the determination of cell-potential in non-standard conditions where the constant temperature of 25oC and molar concentration of 1M is different (LibreText, 2020). Using this principle of the Nernst equation the researcher then calculated the amount of the unknown concentration in the experiment, you can use the equation as shown in equation 3 of the introduction. Thus, according to the presented and calculated data, the zinc solution was dissolve in the reaction and the electron are transferred to the copper ion. Moreover, the amount of the missing concentration value was found according to the use of the Nernst Equation. CONCLUSION In this laboratory, the researcher was able to fully understand the objectives of the experiment. For Table 1.1 of the experiment, the researcher was able to understand the use of cell potential in calculating the voltage produced by the redox couples, in which it can be identified by the movement of the electrons, anions, and cations to determine the cathode and anode of the redox couple. In Table 1.2, we identify the effects of concentration change on cell potential. The change can be identified by the difference in concentration of the solutions used. Such as if we increased the product concentration the cell potential will decreased and vice versa. Lastly, in table 1.3, to get the unknown concentration the used of the Nernst equation was necessary. Since accordingly to its principle, it enables the determination of cell-potential in non-standard

conditions where the constant temperature of 25oC and molar concentration of 1M is different (LibreText, 2020). Additionally, there was a huge discrepancy between the measured and theoretical value of the cell potentials in table 1.1. It maybe because of a systematic error in using the laboratory equipment and an error in calculating the theoretical cell potential. So, it is recommended to double check the equipment’s to be used in the experiment to prevent miscalculation in the amount of the solutions and thus in the solving of the required values. Moreover, manipulating the Nernst equation in table 1.3 to find the unknown concentration of the solution is deeply recommended.

REFERENCES Concentration cell. (n.d.). Retrieved from Khan Academy: https://www.khanacademy.org/science/ap-chemistry/redox-reactions-andelectrochemistry-ap/cell-potentials-under-nonstandard-conditions-tutorialap/v/concentration-cell Electrochemistry. (n.d.). Retrieved from Lumen: https://courses.lumenlearning.com/cheminter/chapter/electrochemistry/ Introduction to electrochemistry. (n.d.). Retrieved from Sparksnotes: https://www.sparknotes.com/chemistry/electrochemistry/intro/terms/ Salt Bridge. (n.d.). Retrieved from Byju's: The Learning App: https://byjus.com/jee/salt-bridge/ Standard reduction potential. (2020, August 16). Retrieved from Libretext: https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_( Analytical_Chemistry)/Electrochemistry/Redox_Chemistry/Standard_Reduction_Potenti al

APPENDIX...