Energy report
a9o2p5
FuelCells.docx
Examining the Voltage of a Hydrogen Fuel Cell with Fuels of Varying Temperatures
CHEM 1B
Jordan Wolz
December 6, 2017
Dates Performed: November 13-27, 2017
Partners: Edwin Chavez, Nhi Nguyen, Jessie Wang, Fletcher Freisen
Instructor: Veronica Jaramillo
1 Introduction
Since the Industrial Revolution, the world has consistently had shortages of energy. The world has been largely dependent on fossil fuels as energy sources. While the Fossil Fuels have been dependable, the fuel is a finite energy source, and eventually the world will not have enough fossil fuels to support our systems. In addition, the negative environmental impact of mankind’s fossil fuel dependency is well-documented. The search for an alternative energy source is alive and well.
Hydrogen Fuel Cells have emerged as a potential solution to the never-ending energy crisis. Hydrogen Fuel Cells generate electricity from the oxidation of a Hydrogen Proton with an oxidizing agent (generally Oxygen from the air). This reaction creates H2O. There are many benefits to using Hydrogen Fuel Cells: they are found to have higher efficiency rates than standard diesel or gas engines, the byproduct (H2O) of the cell is environment friendly, and the fuel is readily available wherever there is water. Unfortunately, platinum is often at the anode to catalyze the reaction that separates the Hydrogen proton from the fuel is expensive and has made it such that no Hydrogen Fuel Cell Developer is yet to turn a profit.
This experiment examines the Voltage produced by a Hydrogen PEM Cells (a PEM cell is a fuel cell that specifically only allows the Hydrogen Proton to pass through the membrane that acts as the ‘salt bridge’ of a galvanic cell. PEM stands for Proton Exchange Membrane) that has Fuel inserted at different temperatures with the goal of better understanding the relationship between the temperature of two reactants of a galvanic cell and the temperature’s affect on the resulting voltage.
2 Procedures
3 Results
Table 1. Collected Data and Selected Calculated Variables
|
Temperature (oC) |
Borax Solution Volume (mL) |
Volume HCl added (mL) |
Tetrobate (M) |
Equilibrium Constant |
lnK |
Inverse Temperature (1/K) |
ΔSo (j/K*mol) |
|
50.0 |
8.8 |
39.2 |
0.45 |
0.36 |
-1.01 |
0.003096 |
299.51 |
|
40.0 |
8.0 |
37.5 |
0.47 |
0.42 |
-0.86 |
0.003195 |
310.61 |
|
30.2 |
8.8 |
22.0 |
0.25 |
0.064 |
-2.74 |
0.003298 |
305.21 |
|
20.0 |
8.8 |
14.9 |
0.17 |
0.020 |
-3.91 |
0.003413 |
306.90 |
|
10.0 |
8.8 |
7.5 |
0.09 |
0.0026 |
-5.97 |
0.003534 |
301.78 |
|
HCl Molarity (M) |
0.202 |
(-1)ΔHo/R (j/K*mol) |
-11962 |
ΔHo (j/mol) |
99452.07 |
Average ΔSo |
304.8 |
|
|
|
|
|
|
|
Standard Dev |
4.35 |
Chart 1. Linearized Graph of Data
Table 2. Percent Error
4 Calculations
Please refer to submitted Calculations.
5 Discussion
This lab has inaccuracies. Using the data from the KHP titration, KHP was determined to have a Ka value of 6.9 x 10-6. The accepted value of KHP is 3.9 x 10-6. This was was accurate enough to calculate the Ka to the correct order, but that is where the compliments end. The Percent error of the observed Ka is over 75%, which is bad. For there to be 0 % error, the pH at the half-equivalence point would be 5.41. The Half-Equivalence Volume is ~9.03 mL. From the data, pH of ~5.41 occurs when about 13mL of NaoH is added.
Despite this, the KHP titration curve seems wonders better than the diprotic titration curve. While the KHP curve leads to interesting Ka values for KHP, the curve at least seems consistent with itself, behaving like a titration curve should with steady increases following a nice clean pattern. Meanwhile, the Diprotic Curve jumps up and down so much, it’s difficult to tell if the first Equivalence Point is an Equivalence Point, or just some weird thing the graph is doing. Rather than steadily increasing, the pH of the diprotic acid would occasionally go down after the base of NaOH was added. This does not seem consistent with the theory that has been taught in class, as NaOH would completely dissolve, with OH- only raising pH and Na+ acting as a spectator ion with no affect on pH.
A plausible explanation for the behavior of this diprotic acid is that the solution was contaminated with some other chemical that would cause the acid in the solution to re-bond with the lost proton. The electrode was also prone to turning off mid-measurement, perhaps causing problems with calibration or consistency with the data.
Aside from these possible error-causers, the experiments were completely caught off guard by the Equivalence Volumes of the diprotic acid when titrating, meaning that as the Equivalence Volumes approached, the experimenters were still adding large amounts of NaOH leading to inaccurate interpolations of where the Equivalence Volume actually happens, and what the pH is at those points.
Despite these errors, it can be said with some degree of confidence that the Unknown Acid is o-phthallic acid. Despite not being totally sure of the exact location or exact pH of the equivalence points and half equivalence point, it does not seem that the data gathered is completely wrong. The data more or less follows what an idea titration curve should look like, and one can tell where the buffer region is centered / about where an equivalence point is, so the pKa’s gathered do have some degree of confidence. When viewing the data, there is an obvious Equivalence Point that sticks out; the second point equivalence point. Using this equivalence point to solve for the molecular weight of the unknown acid gives only a 9 % error when comparing with o-phthallic acid, which is the second best error. On top of that, the known pKas do fall pretty close to where one would expect half-equivalence points to be. Take a gander at the % error table above, and it’s obvious which acid has the lowest % error. For these reasons, it is believed the unknown acid is o-phtallic acid.
In the future, it would behoove the experimenters to make sure to totally clean the equipment prior to use (and grab equipment the experimenter is confident in working), and to lower the titrant amount even before expecting an equivalence point (so as to not miss it) and to make sure to record everything. Practicing these types of behaviors will result in better data, no doubt.
1/T vs lnK (Dissolution of Borax into Tetrobate and Soium ions)
[CELLRANGE] [CELLRANGE] [CELLRANGE] [CELLRANGE] [CELLRANGE] 3.0959752321981426E-3 3.1948881789137379E-3 3.2981530343007917E-3 3.4129692832764505E-3 3.5335689045936395E-3 -1.0098348493659739 -0.85691175141315457 -2.7427377296802771 -3.9117814509009836 -5.9711560281284308 50.0 40.0 30.2 20.0 10.0 Inverse of Temperature (1/T)) Natural Log of the Equilibrium Constant
