Determining the Molar Mass of a Metal by Electrolysis through the Oxidation of Iron, Papers of Chemistry

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Determining the Molar Mass of a Metal by Electrolysis through the

Oxidation of Iron

Lindsey Booth

CCV, Gen Chem II

Tracy Jensen

Abstract It is hypothesized that the molar mass of a metal can be determined by electrolysis using the oxidation of Iron. This is because iron was oxidized at the anode, the positive pole, of the electrolysis cell resulting in the loss of its mass. The slight difference in mass was able to be determined because the initial and final masses of the iron sample were taken before and after electrolysis. For both trials 1 & 2, the difference in initial and final masses was the exact same. Reduction occurred at the cathode, the negative pole, of the electrolysis cell involving the Hydrogen ion. The volume of Hydrogen gas that was produced was measured once 25 mL of gas had been collected in the buret. A meter stick was used to measure the amount of empty space that was left in the buret, from the top of the buret to where the water started. Once this was found, the moles of H gas accumulated could be calculated using the ideal gas law; the number of moles of electrons that passed through the cell was now known. The charge of the cations made led to the determination of the molar mass of Iron. Both trials calculated a higher molar mass than the actual of Iron which is 55.8 g/mol. Trial 1 was highest at 66.3 g/mol, and Trial 2 calculated a molar mass of 63.1 g/mol. Both were not far off, which does provide support for the hypothesis, but experimental error is indicative. Introduction During this lab, electrolysis was studied for two trials using an electrolysis apparatus. Electrolysis was undergone to determine the molar mass of Iron. During electrolysis, the oxidation of the metal Iron occurred, while Hydrogen gas was reduced. Many measurements were taken including the initial and final masses of the Iron sample, before and after electrolysis were taken; the production of Hydrogen gas also provided values of volume, temperature, partial pressure, and height. These measurements contributed to the calculation of molar masses of Iron for each; they were then able to be compared to the actual value.

Dalton's Law states that in a mixture of non-reacting gases, the total pressure exerted is equal to the total of the partial pressures of individual gases^4. The equation is Ptotal = P1 + P2 + P3. Methods Per each trial, an electrolysis apparatus was constructed. A 150 mL beaker was first filled with 0.5 M HC conducting solution about ¾ of the way. Then an iron strip was lightly sanded with steel wool and rinsed with 0.1 M acetic acid, and dried to assure that it was clean. To obtain its initial mass, it was then weighed using a laboratory scale, mass was recorded in the data table. This was done by holding a buret on a clamp stand, and filling it with 50 mL of water completely to the top, allowing no room for bubbles. The buret was then placed and lowered into a 150 mL beaker by placing one finger over the opening and inverting it downward, once again allowing no room for bubbles. An insulated copper wire, serving as the cathode, was then placed into the beaker and up into the opening of the buret. The copper wire cathode was connected to the negative pole of the power supply using an alligator clip. The iron strip anode was placed into the beaker and connected to the positive pole of the power supply using an alligator clip. Both alligator clips, per each electrode, were not submerged into the water within the beaker. Once the electrolysis apparatus was assembled the power supply was turned on, the voltage was slowly increased until the bubbling of hydrogen gas from the cathode was observed. Power was supplied until about 25 mL of H gas was produced, then the volume of Hydrogen was taken. This was done by measuring the empty space between the remaining water and the top of the buret

with a meter stick in millimeters. The temperature of the gas was also recorded using a thermometer and compared to the room temperature. Once those values were collected, the electrolysis apparatus was deconstructed. The iron strip was once again rinsed with water and 0.1 M acetic acid; a paper towel was used to assure that there was no remaining coating of solution on the metal. Afterward, the final mass was taken off the iron strip using a laboratory scale. The conducting solution was placed back into the beaker and used again for the second trial. The electrolysis apparatus was assembled for a second trial, and the same steps were taken. After carrying out the procedures of electrolysis, calculations were performed in order to determine mole pressure, mole values, and molar mass of the Iron sample for each trial. Results Trial 1 Trial 2 Mass of Metal Anode Before Electrolysis (g)

Mass of Metal Anode After Electrolysis (g)

Loss in Mass by Anode (g) 0.055 0. Volume of Gas Collected (mL) 28.47 30. Height of Water in Buret (mm) 353 372

Once the reactions were finished and the experimental part of the laboratory was concluded, calculations were able to be done. The pressure of H2 gas was calculated by plugging pressure values into this equation: Patm = PH2O + PH2 + PWater Column. Using values from the data table, the equation for Trial 1 looked like this: 765.81 = 19.827 + PH2 + 28.460. The pressure of Hydrogen gas was left to solve for, which came out to be 717 mmHg for Trial 1. Trial 2 had an equation which looked like this: 765.81 = 19.827 + PH2 + 30.00, the pressure of H2 came out to be very close at 716 mmHg. After calculating the pressure of hydrogen gas for both trials, the number of moles of Hydrogen gas produced in each trial was able to be calculated, using the ideal gas law. Using PV = nRT. All values except n were plugged in, as we were solving for the number of moles. Trial 1 had 0.01108 moles of hydrogen gas produced, and Trial 2 had 0.01157 moles produced. Once again, both trials produced quite similar values. After the number of moles of Hydrogen gas per trial were calculated, the moles of electrons used to produce the hydrogen gas were able to be calculated. This was done by multiplying each mole value by 2, as there was a 2:1 stoichiometric ratio. For Trial 1, the moles of electrons used to produce the Hydrogen gas was 0.02216 mol, and for Trial 2 the calculated value was 0.02314 mol. The moles of metal oxidized were then calculated for both Trials; this was done by dividing the initial mole values of gas produced by

2. Trial 1 had a value of 0.00554 mol, and Trial 2 had a value of 0.005785. Finally, the molar mass of Iron was able to be calculated per trial. This was done by dividing the difference between the initial and final masses of Iron after and before electrolysis by the number of models of hydrogen gas produced. For Trial 1, 0.55 g was divided by 0.01108, and was calculated out to be 49.6 g/mol. For Trial 2, 0.55 g was divided by 0.1157, and was calculated out to be 47.5 g/mol. The actual molar mass of Iron is 55.8 g mol, so for Trial 1 the difference between calculated and actual was 6.245, and for Trial 2 the difference was 8.3.

Experimental error is inevitable in laboratory experiments, and it is evident that this occurred as there is a significant enough difference between the calculated molar masses and the actual molar mass of Iron. While the difference isn’t drastic, the accumulation of small errors were most likely the contributors. Sources of error in this experiment could’ve most definitely been encountered when measuring the volume of the gas collected as we had to use a meter stick and hold it parallel to the buret, taking our best measurement of the accumulation of gas. Also, measuring the temperature of the Hydrogen gas was tricky as we had to roughly estimate based on the room temperature. Another error could have been caused by insufficient or excessive electrolysis time which could’ve led to incomplete reactions and production of hydrogen gas or Iron oxidation; which could have resulted in different volumes and mass measurements. Conclusion It was hypothesized that the molar mass of a metal could be calculated by electrolysis through the oxidation of Iron. This is supported by calculated molar masses of Iron from Trial 1, 49.6 g/mol, and Trial 47.5 g/mol and comparing these values to the actual which is 55.8 g/mol. While there is a difference, the significance between the calculated and the actual can be attributed to the fact that experimental error exists and was prevalent. References

1. Understanding Half-Reactions , www2.chem.wisc.edu/deptfiles/genchem/netorial/rottosen/tutorial/modules/electrochemistry/02ha lf_reactions/18_21.htm#:~:text=A%20half%2Dreaction%20is%20the,completely%20describe% 0a%20redox%20reaction. Accessed 10 Aug. 2023.