Lab 17 Electrochemistry PDF

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Lab 17: Electrochemistry Name: Xavier Valencia Lab Partner: Trinity Campagna Section: 2A38 4/20/2019 Introduction: A redox equation breaks into two parts, an oxidation and a reduction reaction. Reduction reactions gain electrons, and their charge is reduced. An oxidation reaction loses electrons. The driving force of reaction is an electrochemical potential. For reductions as long as electrons are available, the more readily the species is to be reduced. Electrochemical cells allows for separate redox reactions, where the oxidation end is the anode and the reduction end is the cathode. The overall potential ( Δ E o ) o - Eoanode is the sum of half-cell Δ E o = Ecathode Cell Potential and Gibbs Free Energy ( Δ G ): Δ G=−nF Δ E Cell Potential and Reaction Quotient (Q): RT Δ E = Δ Eo ln(Q) = Δ E o nF

RT ln( nF

[ C]c [ D]d ) [ A ]a [ B ]b Nerst Equation: Δ E = [ Eocathode -

0.0592 o log(Qcathode)] - [ Eanode n

0.0592 log(Qanode)] n

Calomel Half Cell: Hg2Cl2 (s) + 2e- → 2Hg (l) + 2Clo E is the reduction of Calomel Half Cell, and is 0.2458. The sign of Δ E depends on the direction of reaction. In this experiment the Ag electrode is the cathode and the mercury is an anode. In the reaction as well, the silver electrode would be attached to the positive and the mercury electrode would be connected to the negative. The four reaction in this experiment are as follows:

(1) AgI(s): AgI(s) ⇔ Ag+ + I(2) AgCl(s): AgCl(s) ⇔ Ag+ + Cl(3) Ag(NH3)2+:

KAgI = [Ag+][I-]

KAgI = [Ag+][Cl-]

2

N H 3¿ ¿ +¿ N H 3 ¿¿2 KAgI = Ag ¿ ¿ Ag+¿¿ ¿ ¿

Ag(NH3)2+ ⇔ Ag+ + 2NH3

(4) Ag(SCN-)2: −¿ ¿ SCN ¿ 2 Ag ¿ ¿ KAgI = N −¿¿ SC ¿ Ag+¿¿ ¿ ¿

Ag(SCN)2- ⇔ Ag+ + 2SCN-

2

Experimental: Procedure: (found in Green Lab Manual Zhao, General Chemistry Lab Manual (1st Edition), pgs. 206-215) Original Procedure: No procedures were altered Altered Procedure: No procedures were altered Experimental Data: Solution

[Ag+]

1

10-4 M

2

10-3 M

Sample

Δ E

EoAg+/Ag

273.30

0.2733

0.7559

338.10

0.3381

0.7615

Average:

0.7587

Δ E (mV)

Δ E (mV)

Δ E (V)

Calculated [Ag+]

AgCl

-36.2

-0.0362

5.30 x 10-10

Ag(NH3)2+

-105.4

-0.1054

3.60 x 10-11

Ag(SCN)2-

-178.6

-0.1786

2.09 x 10-12

AgI

-394.5

-0.3945

4.70 x 10-16

K eq .

[ions] 0.436 M −¿ ¿ Cl

AgCl

+¿ NH 3 ¿2¿ Ag ¿

0.799M NH 3

−¿ ¿ SCN ¿2 Ag ¿

0.194M −¿ SCN ¿

AgI

0.247M

K accepted

2.31x10-10

1.77 ×10−10 -9.6363

2.298x10-11

1.70 ×10

7.87 x 10-14 −¿ ¿ I

log( K eq. )

1.16 ×10

−16

o

o

0.2733 = [ E Ag+¿ Ag

% error

K accept )

-9.7203

0.86%

-10.541

-6.801

54.99%

1.03 × 10−10 -13..104

-9.99

31.17%

-16.0701

0.839%

−7

−1

8.51 × 10

-15.935

Sample Calculations: E oAg+¿ Ag Calculation: Δ E = [ E Ag +¿ Ag

log(

+¿ ¿ Ag - 0.2458 - 0.0592 * log ¿ ] 1 ¿ 1 - 0.2458 - 0.0592 * log ] −4 [ 10 M ]

E oAg+¿ Ag = 0.7559 [Ag+] Calculation: +¿ ¿ Ag Δ E = [ EoAg +¿ Ag - 0.2458 - 0.0592 * log ¿ ] 1 ¿ +¿¿ Ag −0.0362 = [ 0.7587 - 0.2458 - 0.0592 * log ¿ ] 1 ¿ 1 = 5.30 x 10-10 [Ag+] = 9.275 10❑ Initial [Ag+]:

mL∗1 L ∗.01 mol 1000 mL 1.0 ×10−4 moles =¿ 10.0 1L Initial [Cl ]: mL∗1 L ∗0.50 mol 1000 mL =0.035 moles 70.0 1L Final Concentration [Cl-]: L∗5.30 × 10−10 mol ) = 9.99x10-5 moles Δ Ag = 1.0 ×10−4−(0.10 1L Cl-final mol = (0.035mol) - (9.99x10-5) = 0.0349 mol [Cl ]final = 0.0349 / 0.080 L = 0.436M Keq: Keq = [Ag+][Cl-] Keq = [5.30x10-10][0.436] = 2.31x10-10 log(Keq) = -9.636 Percent Error: −9.636+9.7203 ×100 = 0.86% % Error=¿ −9.7203

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Discussion Questions: 1. All experimental and accepted values can be found in the table above 2. In Oxtoby, Gillis and Nachtrieb a standard hydrogen electrode is used, not like the experiment in which we use a calomel electrode. This change alters the standard reduction potential ¿ −4 3. Δ E=[ Eo cathode −0.0592 log (1.0 ×10 )]−¿ Δ E=[ 0.759 V −0.0592log ( 1.0 × 10−4 )]−[(−0.759 V −0.0592 log (1.0 × 10−2))] Δ E =0.522 V −0.6406 V Δ E =1.163 V 4. If the electrodes were both connected the electrons would discharge too fast to read, almost immediately or too fast to read with the voltmeter. Post Lab Discussion: The purpose of the experiment was to calculate the equilibrium constants for reactions AgCl, Ag(NH3)2+, Ag(SCN)2- and AgI. The calculated equilibrium concentrations of all the reactions are 2.31x10-10, 2.30x10-11, 7.87x10-14 and 1.16x10-16 respectively. The values were then compared to accepted values found in literature. The calculated and literature values were then put on a logarithmic scale in order to determine percent error. The percent error values were calculated to be 0.86%, 54.99%, 31.17% and 0.839%. The high percent error values of the Ag(NH3)+2 and Ag(SCN)2- could derive from the

improper measurement of volumes needed for the experiment, the two low values could indicate that we were precise at a point but made a measurement error for the two previously mentioned experiments. Another possible error can occur from quick back to back solution mixing and not allowing the solution to be thoroughly mixed before measuring the volts and going on to the next step which could lead to high imprecision. In order to improve the accuracy of the next attempt a few things can be changed. For example, double checking that the voltmeter is properly calibrated, or using more precise volumetric measures. For example if the solution needs 15mL use a 25mL graduated cylinder and not a 50 or 100 mL graduated cylinder. Also another huge improvement could be ensuring students thoroughly mix the solution as to fully dissolve and let the equation run to completion. Conclusion: Overall the experiment could be seen as a failure. Although we did collect data that was very close to the expected value for AgCl and AgI (0.86% and 0.839% respectively), the extremely high values of Ag(NH3)+2 and Ag(SCN)2- over shadow the possible successes. Generally students in the lab got a better idea of electrochemistry as a process....