Electrochemistry Report 2019-3 PDF

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Electrochemistry Lab Report...

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Electrochemistry Lab Report Introduction: Electrochemical reactions relate electrical and chemical energy through the combination of redox reactions. In an electrochemical cell, the reduction half-reaction and the oxidation half-reaction are split up in space. Species are reduced at the cathode and species are oxidized at the anode. To determine the overall potential of the cell, you can use the following equation: E°overall = E°cathode- E°anode The Nernst equation is important in the study of electrochemistry because it relates electrochemical reactions, Gibbs Free Energy, and chemical equilibrium. The Nernst equation can take the following form: ΔE = [E°cathode - (0.0592/n)(logQcathode)] - [E°anode - (0.0592/n)(logQanode)]; where n is the number of moles of electrons and Q is the reaction quotient. In this lab, using the calomel half-cell as the reference electrode, the standard electrochemical potential for a silver electrode was determined experimentally. Once this is accomplished, the calomel electrode and the silver electrode are used to measure the silver ion concentration for different equilibrium reactions of silver compounds. The following equation is helpful in determining the silver ion concentration: ΔE = [E°Ag+/Ag - .2458 - (.0592)(log(1/[Ag+ ]))] This equation is derived from the Nernst equation; the value of .2458 is the whole reduction potential from the calomel half-cell, including the effect from the [Cl- ]. This lab provides the opportunity to understand the concepts and set up of electrochemical reactions. Understanding electrochemistry is important for its practical applications. Electrochemical cells are used to purify metals produced by the processes of pyrometallurgy. In addition, they are used in the development of batteries. We will be exploring the following reactions and experimentally determining their equilibrium constants:

Reaction 1: AgI(s) Ag+ (aq) + I- (aq); KAgI = [Ag+ ][I-] Reaction 2: AgCl(s) Ag+ (aq) + Cl- (aq); KAgCl(s) = [Ag+ ][Cl-]

Reaction 3: Ag(NH3)2 + Ag+ (aq) + 2NH3(aq); KAg(NH3)2+(s) = [Ag+][NH3]2/[Ag(NH3)2 +] Reaction 4: Ag(SCN)2 - Ag+ (aq) + 2SCN-(aq); KAg(SCN)2-(s)= [Ag+][SCN-]2/[Ag(SCN)2 -] The calomel half-cell will serve as our known electrode; this standard reduction potential will be essential for all calculations. The reaction in the calomel cell is as follows: Hg2Cl2(s) + 2e- --> 2Hg(l) + 2Cl- Experimental:  For this experiment, one must first create a 10-4  and 10-3  dilution of AgNO3 from stock solution; the most precise method of doing this is through serial dilutions. After creating these solutions, the experimenter should measure the cell potential for these dilutions by filling the test tube with each solution and leveling the liquid height with the calomel cell. Insert the silver electrode into solution; attaching the red clip of the voltmeter to the silver and the black clip of the voltmeter to the copper of the calomel cell will set up your cell. Record the voltage. Place 100-mL of water into a beaker and place on a stirrer. First, add 10.0 mL of 0.01 M AgNO3 and 70.0 mL of 0.50 KCl solution to attain the AgCl equilibrium. Record the voltage after attaching the red clip of the voltmeter to the silver and the black clip to the mercury. Add 20.0 mL of 4.0M ammonia to the same solution to attain the Ag(NH3)+ equilibrium. Record the polarity and the voltage. Add 1.9 g of KSCN to the same beaker. Wait for stabilization and record the polarity and voltage. Add 3.3 g of KI to the same beaker. Wait for stabilization and record the polarity and voltage. Record qualitative observations at each step and dispose of each solution in the appropriate waste bin. [No deviations were made from the UChicago General Chemistry Lab Manual] Data Analysis: Qualitative During the determination of the standard potential of the silver electrode, the polarity was positive. During the equilibrium reactions portion, the cell potentials all had a negative polarity. For all of these reactions, the black clip was on the silver electrode and the red clip was on the mercury electrode.

In the AgCl equilibrium, the solution was a murky white because of the presence of a precipitate (presumably KNO3). The addition of the ammonia to the AgCl mixture produced a strong smell. The solution also transitioned from murky to clear. The addition of the white, snow-like crystals of KSCN(s) didn’t produce any physical changes; the solution remained clear. The addition of KI(s), a rocky white solid, changed the solution to a murky yellow. Quantitative Table 1: Standard Potentials Solution #

[Ag+]

Potential ΔE (V)

E°Ag+/Ag (V)

1

10-4

.258

.741

2

10-3

.345

.887

Mean E°Ag+/Ag = 0.814 Sample Calculation for E°Ag+/Ag, Solution 1: ΔE = [E°Ag+/Ag - .2458 - (.0592)(log(1/[Ag+ ]))] 0.258V = [E°Ag+/Ag - 0.2458 - (.0592)(log(1/[10-4  ]))] 0.258V = [E°Ag+/Ag  - 0.2458 - .2368] E°Ag+/Ag = .7406 V Table 2: Silver Concentrations in Each Sample Sample

ΔE (V)

Calculated [Ag+ ] (M)

AgCl(s)

-.029

8.26E-11

Ag(NH3)2+(aq)

-.104

4.47E-12

Ag(SCN)2-(aq)

-.173

3.21E-13

AgI(s)

-.387

7.41E-17

Sample Calculation for [Ag+ ], AgCl(s): ΔE = [E°Ag+/Ag - .2458 - (.0592)(log(1/[Ag+ ]))] -.029 = [.8137 - .2458 - (.0592)(log(1/[Ag+ ]))] 10.1 = log(1/[Ag+ ]) [Ag+] = 8.26E-11 Table 3: Anion Concentrations in Each Sample Sample

Species

Calculated Concentration of

Calculated [Ag+ ] (M)

Equilibrium Constant

Species (M) AgCl(s)

[Cl- ]f

0.436

8.26E-11

3.60E-11

Ag(NH3)2+(aq)

[NH3 +]f

0.798

4.47E-12

2.85E-9

Ag(SCN)2-(aq)

[SCN- ]f

0.198

3.21E-13

1.26E-11

AgI(s)

[I- ]f

0.198

7.41E-17

1.47E-17

Sample Calculation for [Cl- ]f, AgCl(s): [Cl- ]f = ((0.5M)(0.07L)/(0.08L)-((0.01M)(0.01L)/(.08L)) = 0.436M Sample Calculation for Equilibrium Concentration, AgCl(s): K = [Ag+ ]f[Cl- ]f K = (8.26E-11)(0.436) K= 3.64E-9   Note: for repeating this calculation for Ag(NH 3 ) 2 + (aq) and Ag(SCN) 2 - (aq), the coefficient for the anion   in the equilibrium concentration is 2; the term is squared. In addition, the calculation for the moles of solute is made easier by simply taking the mass in grams and dividing by the molar mass (for Ag(SCN) 2 -  and AgI).

Table 4: Error in Equilibrium Constants Sample

Equilibrium Constant

Literature Value of Equilibrium Constant

Logarithmic Percent Error

AgCl(s)

3.60E-11

1.60E-10

6.61%

Ag(NH3)2+(aq)

2.85E-9

5.88E-8

18.2%

Ag(SCN)2-(aq)

1.26E-11

unknown

unknown

AgI(s)

1.47E-17

1.50E-16

6.37%

Sample Calculation for Logarithmic Percent Error, AgCl(s): |log(observed)-log(literature)|/log(literature) = decimal error*100 = percent error |(log(3.60E-11)-log(1.60E-10))|/log(1.60E-10)| = 0.0661 = 6.61%

Discussion: 1. In this lab, the standard reduction potential for silver was calculated by taking the average of two experimentally calculated potentials. The experiment produced an average value of 0.814 V. Compared to the literature value of .7996 V, there is 1.77% error. In the second portion of the lab, using the cell potential determined in the experiment, the concentrations of the silver ions in different equilibrium reactions were calculated using the Nernst equation and can be found in Table 2. The final concentrations of the anions in each compound were also calculated and compiled in Table 3. Knowledge of both of these concentrations allowed for the calculation of the equilibrium constants of 3.60E-11 for AgCl, 2.85E-9 for Ag(NH3)2 +, 1.26E-11 for Ag(SCN)2 -, and 1.47E-17 for the AgI. The experimental and literature values can be found in Table 4. All of our values were smaller than literature value. With these experimental values the percent error of log(Keq  ) was calculated to be 6.61% for AgCl, 18.2% for Ag(NH3 )2 +, and 6.37% for AgI. The literature value for Ag(SCN)2 - could not be found so a percent error was not calculated. All literature values were obtained from Oxtoby's Principles of Modern Chemistry. Because of these percent errors, it is necessary to examine potential sources of errors. Several factors contribute to this measurement error, including the contribution to the potential from interfering ions, the variations in temperature, and the sensitivity of the voltmeter. The presence of other ions that are not desired can affect the measurement of the cell potential that in turn affects the calculation of the equilibrium constant. Because the calomel reference electrode is dependent on the concentration of the chloride, and the concentration of the chloride is dependent on the temperature, variations in temperature can produce inaccurate potentials. Finally, the sensitivity of the voltmeter can dramatically change the way the current passes through the wire and have a negative effect on accuracy. All these sources of are very difficult to control. Making sure the solutions are pure and the temperature is constant are two ways to resolve these issues. 2. Although the standard reduction potential of the calomel half-cell is .268 V, the value used in the experiment was .2458 volts because this value consists of the whole reduction potential including the contribution from the concentration of the chloride and the 2 moles of electrons generated in the anode. 3. If a half –cell with 10-2  M AgNO3(aq) replaced the calomel electrode, the cell potential would be -0.1184V. This is calculated using the Nernst equation: [0.814V - (.0592V/1)(log(1/10^-4))] - [0.814V - (.0592V/1)(log(1/10^-2))] = -0.1184V. 4. If the two electrodes were connected directly, there would be no resistance to the flow of electric current. This setup would produce a short circuit. Short circuits can be dangerous with high voltage power sources because they produce overheating and uncontained electricity. Conclusion: This lab provided an opportunity to set up and understand the practical applications of electrochemical reactions. It also provided an understanding of the relationship between electrochemistry and chemical equilibrium through the Nernst equation. In this lab, the silver

standard potential was determined to be 0.6840 V. The experimentally determined potential had an 1.77% error. Using the Nernst equation, the silver standard potential was used to calculate the concentration of silver ions in several equilibrium reactions involving silver compounds. With this knowledge the equilibrium constants for each reaction were calculated: 3.60E-11 for AgCl, 2.85E-9 for Ag(NH3)2 +, 1.26E-11 for Ag(SCN)2 -, and 1.47E-17 for the AgI. The percent error of  log(Keq) was calculated to be 6.61% for AgCl, 18.2% for Ag(NH3)2 +, and 6.37% for AgI. The literature value for Ag(SCN)2 - could not be found so a percent error was not calculated. Possible  sources of error included interfering ions, variations in temperature, and the sensitivity of the voltmeter. Using pure solutions and keeping the temperature constant can minimize some of the error. It is important to not connect the electrodes directly because it can create a short circuit and generate unwanted heat energy. Sources Cited: 1. Zhao, Meishan and Dragisich, Vera. General Chemistry Experiments . Hayden-McNeil Macmillan Learning, 2018. 2. “Solubility Table.” Wikipedia , Wikimedia Foundation, 13 Mar. 2019, en.wikipedia.org/wiki/Solubility_table#cite_ref-2....