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To determine the standard molar volume of a gas at a specific pressure and temperature, using calculations involving gas laws.
The controlled variables in this lab are the purity of the materials used.
Magnesium must be as pure as possible, limiting oxidization.
Hydrochloric acid must be kept at the same concentration for the lab to be reproduced, and ideally come from the same source.
Water must come from the same source (and ideally, distilled). In this lab, the water came from the water tap.
The independent variables are the mass of the magnesium used and the volume of the hydrochloric acid used, since these are the reactants used in the experiment.
The dependent variable is the hydrogen gas produced, since it is the product of the reaction between the magnesium and the hydrochloric acid.
Silver coloured.
Firstly, the moles of hydrogen gas produced must be calculated using the mass of magnesium and the balanced chemical equation of the reaction.
The balanced chemical reaction: Using this balanced chemical reaction, the moles of hydrogen can be produced using the mass of magnesium.
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Uncertainties must also be calculated. In this case, the only uncertainty lies with 0.03 ± 0.01g Mg.
Secondly, the pressure of the hydrogen gas collected must be calculated, using the total pressure and the pressure of water vapor at the temperature. This can be calculated using Dalton’s Law of partial pressures.
Thirdly, the volume that the gas occupies at STP must be calculated using the observed volume and the corrected partial pressure of hydrogen using the combined gas law: P2, T2 and V2 represent the variables at STP.
All temperatures must also be changed to Kelvin 16 ± 0.5⁰C + 273 = 289 ± 0.5 K
Now the combined gas law can be used. Fourthly, the experimental molar volume of hydrogen gas at STP must be calculated based on the theoretical yield.
Fifthly, the percent error must be calculated Sixthly, the moles of hydrogen gas collected, based on observed temperature, volume and corrected partial pressure must be calculated using the ideal gas law.
Finally, the percent yield of the experiment must be calculated.
To conclude, the moles of hydrogen gas collected were 0.00175 ± 0.00006 mol H2. The theoretical calculation of the moles of hydrogen, however, was lower, at 0.001 mol H2. The pressure of hydrogen, accounting for water temperature, was 99.58 ± 2 kPa. The volume that the gas occupies at STP was 0.0398 ± 0.002 L, while the experimental molar volume at STP was 30 ± 20 L/mol. Thus, the percent error was 30% ± 50%, and the percentage yield was 175% ± 50%.
Using this information, it can be said that the experimental molar volume of hydrogen is 22.7 L/mol at STP, due to the 20L/mol uncertainty on the calculated experimental molar volume. However, the massive uncertainties present are too inaccurate (At 50% for the molar volume) and while the molar volume does technically fit our answer, it is too inaccurate to validate it. To summarise, there is a lack of evidence to prove that the molar volume is 22.7 L/mol at STP, and this lack of evidence is too large to determine the correct value.
Other factors also show that the experiment was not successful; the calculated percent yield is 175% ± 50%, which is theoretically impossible in this experiment; the percent error also has a massive uncertainty, and can technically be -20% accounting uncertainties, which is also impossible. These values thus come from sources of error which have been encountered in the lab.
As seen by the data above, sources of error are clearly present in this lab. Firstly, the use of a small piece of magnesium (with a tiny mass) gave a massive uncertainty of 40% before even attempting the lab, which is far too big to get an accurate result. This systematic error greatly impacted percentage yield, percent error, experimental molar mass and the experimental number of moles of hydrogen. Its lack of significant figures may have also resulted in several rounding errors, which would have resulted in more inaccurate data.
The purity of materials used, especially the concentration of hydrochloric acid, may have also severely impacted the lab, resulting in another systematic error. If the magnesium was coated in magnesium oxide, the reaction may have made less gas than intended, or another unwanted side reaction may have occurred.
A random error which may have impacted the lab was the method of sealing the eudiometer. This method may have caused air to enter the eudiometer and caused a greater overall volume of gas. This would also explain the bigger experimental number of moles compared to the theoretical calculation.
To improve the lab, several things could be done to eliminate the sources of error above. Improving the lab in this way would result in smaller uncertainties and would overall give a more accurate answer.
Firstly, a bigger piece of magnesium (along with a larger volume of hydrochloric acid) must be used. This would eliminate the large uncertainties present before the start of the lab, as well as making a side reaction less impactful on the overall volume of hydrogen created.
Secondly, elements which are as pure as possible should be used as often as possible; this includes using distilled water instead of tap water, using pure magnesium and using hydrochloric acid which is at 2M. Using more pure elements would decrease the risk of possible side reactions resulting in more gas.
Thirdly, a better sealing mechanism should be used when filling the eudiometers with water and hydrochloric acid, as excess air adds to the volume of water. Lightly tapping the side of the eudiometer to rid it of any gas before the reaction would also help with accuracy.
The implication of these improvements would lower uncertainty, increase accuracy and limit random errors present in the lab, which should make it possible to reproduce and obtain and calculate the correct number of moles of hydrogen based on the lab data.
The Standard Molar Volume of a Gas Lab. (2024, Feb 06). Retrieved from https://studymoose.com/document/the-standard-molar-volume-of-a-gas-lab
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