————————- ————————————————- ————————————————- Abstract: ————————————————- ————————————————-
The aim of this lab was to determine the volume of one mole of hydrogen gas at STP (standard temperature and pressure). In the experiment, a magnesium ribbon was placed at the bottom of an eudiometer tube and allowed to react with 2M hydrochloric acid. The volume of hydrochloric displacement was measured and calculations determined the molar volume of hydrogen gas produced inside the eudiometer tube. After performing and taking the average data of two trials, it was concluded that the calculated value of one mole of hydrogen gas at STP was about 22.
35 L/mol. Rounding up, it can be concluded that one mole equals 22. L/mol, which is the accepted value of the volume one mole of an ideal gas takes up. However, there was a 0. 357% error, which could have been due to minor problems such as reading the values on the instruments slightly off. In the end, the experiment demonstrated that one mole of hydrogen gas at STP takes up 22.
4 L. ————————————————- ————————————————- ————————————————- Introduction: An ideal gas can be defined as a hypothetical gas whose molecules occupy negligible space and have no interactions, therefore exactly obeying the gas laws.
However, most of the gases encountered in the real world are real gases which do not strictly obey ideal gas laws. The purpose of this lab is to find the molar volume of H? gas at STP. In this experiment, a known mass of magnesium was reacted with a solution of hydrochloric acid and water in an eudiometer tube. The volume of displacement of the reacted solution would determine the total volume of H? gas and water vapor produced from the reaction. After the data is calculated and corrected for the differences in temperature and pressure, the result is the volume of one mole of hydrogen gas at STP.
It was hypothesized that the calculated molar volume of H? gas from this experiment would be close to the literature value of 22. 43 L/mol, although it would be slightly greater/less than 22. 43 L/mol because of possible sources of error due to inaccurate measurements of the H? gas. The independent variable is the amount of magnesium used for the reaction, and the dependent variable is the volume of H? gas produced. Temperature, pressure, and total volume are assumed to be controlled variables during this experiment. ————————————————- ————————————————- ———————————————— ————————————————- Experimental Procedures: ————————————————- Materials: ————————————————- Hydrochloric acid, HCL, 2 M, 30 mL ————————————————- Distilled water ————————————————- Magnesium ribbon, Mg, 2 ————————————————- Barometer ————————————————- Beaker, 400-mL ————————————————-
Copper wire, Cu, 18 gauge, 15-cm long ————————————————- Eudiometer tube, 50-mL ————————————————- Graduated cylinder, 500-mL, 25-mL ————————————————- Metric ruler ————————————————- One-hole rubber stopper ————————————————- Thermometer ————————————————- Wash bottle ————————————————- Ring stand ————————————————- Test tube clamp ———————————————— ————————————————- Procedure: 1. ————————————————- Fill the 400-mL beaker about ? full with distilled water. 2. ————————————————- Measure the piece of magnesium ribbon. 3. ————————————————- Calculate the mass of the magnesium ribbon with the conversion factor 0. 01468 g/cm. 4. ————————————————- Twist and fold one end of the 15-cm copper wire and insert the magnesium ribbon through it.
Bend the magnesium ribbon around the copper wire to make sure it will not escape during the experiment. 5. ————————————————- Put 15-mL of 2 M hydrochloric acid, HCl, in a 25-mL graduated cylinder. 6. ————————————————- Slowly add the 2 M hydrochloric acid, HCl, to the eudiometer tube while slightly tipping it. 7. ————————————————- Fill the eudiometer tube all the way to the top with water. Avoid mixing the acid with the water. 8. ————————————————-
Insert the magnesium-copper wire-stopper into the tube. 9. ————————————————- Put finger over hole of the rubber stopper, invert the tube and put it into the 400-mL beaker. Clamp the eudiometer tube in place with a test tube clamp and ring stand. 10. ————————————————- Record observations of chemical reactions taking place. 11. ————————————————- If the magnesium ribbon “escapes” the copper wire, shake the eudiometer up and down without lifting it out of the water. 12. ————————————————-
Wait about 5 minutes or when the magnesium ribbon is completely dissolved. Tap the sides to dislodge the gas bubbles. 13. ————————————————- Fill a 500-mL graduated cylinder with tap water. Place it in the sink. 14. ————————————————- Move eudiometer tube up and down in the cylinder until the water level in the eudiometer is the same as the water level in the graduated cylinder. 15. ————————————————- Record the volume of hydrogen gas in the tube. 16. ————————————————-
Record the temperature of the water bath in the 500-mL graduated cylinder. 17. ————————————————- Measure the barometric pressure in the lab. 18. ————————————————- Discard the solution in the sink. 19. ————————————————- Repeat process with second strip of magnesium ribbon. ————————————————- ————————————————- Precautions: * ————————————————- Hydrochloric acid is a corrosive liquid! Avoid direct contact with it. ————————————————- Magnesium metal is flammable. Avoid flames and other sources of ignition. * ————————————————- Wear goggles. (For extra protection, wear gloves and an apron. ) * ————————————————- Wash hands thoroughly with soap and water before and after the lab. ————————————————- ————————————————- Pitfalls: * ————————————————- Possibly inaccurate readings of the volumes, temperatures and pressure. ————————————————- The magnesium ribbon broke into pieces and either floated up or down the eudiometer tube. * ————————————————- The hydrogen gas escaped through the rubber stopper. * ————————————————- Difference in volume due to transferring of the eudiometer tube. ————————————————- ————————————————- Observations/Data: ————————————————- ————————————————- Collected Data ———————————————— | ————————————————- Trial 1| ————————————————- Trial 2| ————————————————- Length of Mg Ribbon| ————————————————- 2. 61 cm| ————————————————- 2. 58 cm| ————————————————- Mass of Mg| ————————————————- 0. 0383 grams| ————————————————- 0. 0379 grams| ————————————————-
Conversion Factor| ————————————————- 0. 01468 g/cm| ————————————————- 0. 01468 g/ cm| ————————————————- Evidence of Chemical Reaction| ————————————————- Bubbles and gas formed when HCl came in contact with the Mg ribbon. | ————————————————- Bubbles and gas formed when HCl came in contact with the Mg ribbon. | ————————————————- Volume of H2 Gas| ————————————————- 39. 1 mL| ————————————————- 39. 50 mL| ————————————————- Corrected Volume of H2| ————————————————- 35. 2 mL| ————————————————- 35. 0 mL| ————————————————- Barometric Pressure| ————————————————- 754 mm Hg| ————————————————- 754 mm Hg| ————————————————- Temperature of Water in Graduated Cylinder| ————————————————- 22. °C| ————————————————- 23. 8 °C| ————————————————- Observations| ————————————————- Since the density of HCl is higher than the distilled water’s density, the HCl sunk to the rubber stopper when the eudiometer was inverted. The HCl started to react with the Mg, and bubbles and gas were formed. The Mg started to dissolve and turned white. After a short while, the solution began decreasing rapidly. A small part of the Mg ribbon broke off and floated to the top of the solution. Some gas escaped from the hole in the rubber stopper.
The formation of bubbles stopped when all of the Mg ribbon was dissolved. | ————————————————- When the eudiometer was inverted, the HCl sunk to the rubber stopper, because it is more dense than distilled water. After approximately thirty seconds, bubbles and gas formed. This was probably because the HCl began reacting with Mg. At around two minutes and thirty seconds, the solution started to decrease quickly. It took over six minutes for the reaction to go to completion. The Mg ribbon dissolved completely with no breakage. | ————————————————- ———————————————— ————————————————- Data Analysis/Graphs: ————————————————- Calculated Results For Each Trial ————————————————- | ————————————————- Trial 1| ————————————————- Trial 2| ————————————————- Theoretical moles of H2 Gas| ————————————————- 0. 00158 mol| ————————————————- 0. 00156 mol| ————————————————-
Partial Pressure of H2 Gas| ————————————————- 733 mm Hg| ————————————————- 732 mm Hg| ————————————————- Corrected Volume of H2 Gas| ————————————————- 0. 0352 L| ————————————————- 0. 0350 L| ————————————————- Molar Volume of H2 Gas| ————————————————- 22. 3 L/ mol| ————————————————- 22. 4 L/ mol| ————————————————- ———————————————— The results for Trial 2 correlated more with the concept of molar volume of an ideal gas. According to the molar volume, for every mole of gas, 22. 4 liters should be filled. This applies to the case of Trial 2 where the molar volume of hydrogen gas was 22. 4 L/mol, while in Trial 1 there was some error, for the molar volume of hydrogen gas was only 22. 3 L/mol. ————————————————- Calculated Results From Both Trials ————————————————- Average Molar Volume of H2 Gas | ————————————————- 22. L/mol| ————————————————- Molar Volume of H2 Literature Value| ————————————————- 22. 43 L/mol| ————————————————- Molar Volume of H2 Gas Percent Error| ————————————————- 0. 357 %| ————————————————- Experimental Density of Hydrogen| ————————————————- 0. 0904 g/ L| ————————————————- Density of Hydrogen Literature Value| ————————————————- 0. 0899 g/ L| ———————————————— ————————————————- The experimental density of hydrogen is greater that the literature value density. This shows that there were more grams of hydrogen per liter in our experiment than in the ideal gas. This error was most likely caused by rounded numbers and limited significant figures. ————————————————- ————————————————- ————————————————- 1. Calculate the theoretical number of mole of hydrogen gas produced in Trials 1 and 2. ———————————————— convert grams Mg to moles Mg and use the molar ratio of Mg and H2 to get moles H2 ————————————————- ————————————————- Trial 1:0. 0383 g Mg 1 mol Mg24. 30 g Mg1 mol H21 mol Mg=0. 001576132? 0. 00158 mol H2 ————————————————- ————————————————- Trial 2:0. 0379 g Mg 1 mol Mg24. 30 g Mg1 mol H21 mol Mg=0. 001559671? 0. 00156 mol H2 ————————————————- ————————————————- ———————————————— 2. Calculate the partial pressure of hydrogen gas produced in Trials 1 and 2. (Refer to the Table of Vapor of Water at Different Temperature) ————————————————- Ptotal =PH2+ PH2O ————————————————- ————————————————- Trial 1 : PH2 = 754 mm Hg – 21. 1 mm Hg = 732. 9 ? 733 mm Hg ————————————————- Trial 2 : PH2 = 754 mm Hg – 22. 4 mm Hg = 731. 6 ? 732 mm Hg ————————————————- ———————————————— ————————————————- 3. Use the combined gas law to convert the measured volume of hydrogen to the “ideal” volume the hydrogen gas would occupy at STP for Trials 1 and 2. ————————————————- convert mm Hg to atm, Celsius to Kelvins, and mL to L, then substitute into P? V? T1= P? V? T? ————————————————- Trial 1 : 733 mm Hg 1 atm760 mm Hg= 0. 964473684 ? 0. 964 atm ————————————————- 22. 8? +273. 15=295. 95? 296 K ————————————————- 9. 61 mL 1 L 1000 mL= 0. 03961 L ————————————————- (0. 964 atm)(0. 03961 L) 296 K= (1 atm) V? 273 K ? V2=273 K0. 964 atm0. 03961 L1 atm 296 K=0. 035217037? 0. 0352 L H? ————————————————- ————————————————- ————————————————- Trial 2 : 732 mm Hg 1 atm760 mm Hg= 0. 963157895 ? 0. 963 atm ————————————————- 23. 8? +273. 15=296. 95? 297 K ————————————————- 39. 50 mL 1 L 1000 mL= 0. 03950 L ————————————————- 0. 964 atm)(0. 03950 L) 296 K= (1 atm) V? 273 K ? V2=273 K0. 964 atm0. 03950 L1 atm 296 K=0. 034964682? 0. 0350 L H? ————————————————- ————————————————- 4. Divide the volume of hydrogen gas at STP by the theoretical number of moles of hydrogen to calculate the molar volume of hydrogen for Trials 1 and 2. ————————————————- ————————————————- corrected volume of hydrogen gastheoretical moles of hydrogen gas = molar volume of hydrogen gas ————————————————- ———————————————— Trial 1 : 0. 0352 L H? 0. 00158 mol H? = 22. 27848101 ? 22. 3 L/ mol H? ————————————————- Trial 2 : 0. 0350 L H? 0. 00156 mol H? = 22. 43589744 ? 22. 4 L/ mol H? ————————————————- ————————————————- 5. What is the average value of the molar volume of hydrogen? Calculate the percent error in your experimental determination of the molar volume of hydrogen. ————————————————- The average value for the molar volume of hydrogen is 22. L/ mol. ————————————————- ————————————————- Calculating the average value of the molar volume of hydrogen: ————————————————- the sum of the molar volumes of hydrogen from the experiment divided by the number of trials ————————————————- (22. 3 L/mol H? + 22. 4 L/mol H? )2 =44. 7 L/mol H? 2= 22. 35? ? 22. 4 L/mol H? ————————————————- ————————————————- 122. 35 L/mol will be used in later calculations for precision. ———————————————— ————————————————- Calculating the percent error of the molar volume of hydrogen: ————————————————- Percent error= |Experimental value-Literature value|Literature value ? 100% ————————————————- | 22. 35 L/ mol H? – 22. 43 L/ mol H? |22. 43 L/ mol H? ? 100% = |-0. 08 L/ mol H? |22. 43 L/ mol H?? 100% = 0. 08 L/ mol H? 22. 43 L/ mol H?? 100% ————————————————- ————————————————- =0. 003566652 ? 100% ? 0. 357% ———————————————— ————————————————- 6. Use your value of the molar volume of hydrogen to calculate the mass of one liter of hydrogen gas at STP. How does this experimental value of the density compare with the literature value? ————————————————- This experimental value of the density of hydrogen gas is slightly greater than the literature value. The literature value is 0. 0899 g/L, while the experimental value is 0. 0904 g/L. The reason for the slight difference is due to the precision of the values used to calculate the density.
For the literature value, 2. 0158 g/mol (molar mass of H2) and 22. 43 L (molar volume of an ideal gas) were used to calculate the density of hydrogen gas. On the other hand, 2. 02 g/mol (rounded molar mass of H2) and 22. 35 L (experimental molar volume of H2) were used to calculate the density of hydrogen gas. ————————————————- ————————————————- Calculating the mass of one liter of hydrogen gas at STP: ————————————————- molar mass H2experimental molar volume of H2=Density of H2 ————————————————- . 02 gmol H222. 35Lmol H2=0. 090380313? 0. 0904g/L H2 ————————————————- ————————————————- 7. In setting up this experiment, a student noticed that a bubble of air leaked into the eudiometer when it was inverted in the water bath. What effect would this have on the measured volume of hydrogen gas? Would the calculated volume of hydrogen be too high or too low as a result of this error? Explain. ————————————————- The bubble of air that leaked into the eudiometer will increase the measured volume of hydrogen gas.
The calculated volume of hydrogen would be too high, because the volume of the air bubble would be included in the measured volume of gas which would in turn affect the calculations of the corrected volume of hydrogen gas. ————————————————- ————————————————- 8. A student noticed that the magnesium ribbon appeared to be oxidized- the metal surface was black and dull rather than silver and shiny. What effect would this error have on the measured volume of hydrogen gas? Would the calculated molar volume of hydrogen be too high or too low as a result of this error? Explain. ———————————————— Since the magnesium ribbon is already oxidized, it would have given up electrons. HCl is usually reduced by the electrons of the magnesium ribbon, but since the magnesium ribbon has been oxidized, there will be less electrons to react with the HCl. Since the electrons reaction with the HCl is what forms the hydrogen gas, the oxidized magnesium ribbon will have less electrons to react with the HCl. Therefore there will be a smaller volume of hydrogen gas produced. The calculated molar volume of hydrogen would in turn be too low. ————————————————- ———————————————— Other Calculations: ————————————————- Calculating the mass of Mg ribbon ————————————————- length of Mg ribbon(cm) ? conversion factor (g/cm) = mass Mg ————————————————- Trial 2:2. 61cm Mg 0. 01468 g Mg1 cm=0. 0383148? 0. 0383 g Mg ————————————————- Trial 2:2. 58cm Mg 0. 01468 g Mg1 cm=0. 0378744? 0. 0379 g Mg ————————————————- ————————————————- Discussion: ———————————————— What can you claim from your results and what evidence leads you to make your claim? ————————————————- With the molar volume of gas, the amount of gas needed to fill any volume could be calculated. The concept of molar volume is 1 mole of a gas occupies 22. 4 liters at STP. In this lab, by finding the volume of hydrogen gas produced under laboratory conditions and then using gas law formulas to calculate the volume that one mole of hydrogen would occupy at STP, the accuracy of the concept of molar volume could be determined. ———————————————— Magnesium metal and hydrochloric acid react at a one to one mole ratio. The conversion factor of grams per centimeter and molar mass was used to determine the moles of magnesium, 0. 00158 moles and 0. 00156 moles. This meant that 0. 00158 moles and 0. 00156 moles of hydrogen gas will be produced. After reacting the magnesium metal with the hydrochloric acid, 39. 61 mL and 39. 50 mL of gas was produced. Using the vapor pressure of water at different temperatures, the partial pressure of the water vapor was calculated.
It was then subtracted from the laboratory pressure, 745 mm Hg, to get the partial pressure of the hydrogen gas. Their pressures were 733 mm Hg and 732 mm Hg. The combined gas law was then used to convert the measured volume to the “ideal” volume. The results showed the volume would have been 35. 2 mL and 35. 0 mL if hydrogen was an ideal gas. To calculate the molar volume, the “ideal” volume is divided by the moles of hydrogen gas that was present. The molar volume came out to be 22. 3 liters per mole of hydrogen and 22. 4 liters per mole of hydrogen gas. The concept of molar volume states that at tandard pressure and temperature (STP), 1 mole of gas would occupy 22. 4 L. The data proves that this concept is correct. The little bit of error from trial one could be caused by the hydrogen gas that was observed escaping the rubber stopper because the magnesium ribbon had broken and fallen right on top of the hole in the rubber stopper. ————————————————- ————————————————- Conclusion: This experiment taught how to collect gas from a reaction in a tube and determine its molar volume given the volume measurement, temperature, and theoretical moles of the gas.
It also shows how real gases differ from ideal gases, so the accepted value for ideal gases of 22. 42 L/mol is close but not a precise value for measuring real gases. It applies to what is being taught in class about molar volume, reading glassware, pressure, and gas laws. In real life, this information could be useful in predicting the behaviors of real gases. Knowledge of gas laws can be applied to opening a bottle/can of soda or aerosol cans. When a can of soda is opened and left sitting, the soda is no longer “fizzy” because the carbon dioxide in the drink escaped to equalize the pressures in the can and in the atmosphere.
Gas laws explains why aerosol cans explode when it gets too hot. The increase in temperature causes an increase in pressure, and too much pressure inside the can causes the can the burst. The molar volume of an ideal H? gas is 22. 43 L/mol, so it was hypothesized that the results would yield a value close to but less than the ideal gas molar volume. The results would be slightly less than the ideal volume because of the possible errors that would arise in experiments with real gases. During the reaction, the hydrogen gas bubbles rose to the top of the eudiometer tube because it has a lower density than water.
However, some of the gas bubbles were too small and were trapped in the liquid. Those trapped H2 bubbles did not contribute to the volume of H2 recorded, so the actual volume of H2 should have been greater. Hydrogen gas is also slightly soluble in water, which leads to a smaller measured volume of H2 gas. Some hydrogen molecules also could have escaped from the small hole at the bottom of the container, because it was observed that there were bubbles coming from the opening that was placed inside the beaker of water.
There is also a possibility of inaccurate volume measurements because the pressure inside the eudiometer could have been not completely equalized with the atmospheric pressure. The calculated average molar volume of H2 gas is 22. 35 L/mol, which is only a 0. 357% error from the accepted volume of one mole of an ideal H? gas, 22. 43 L/mol. This minor error in calculation could be due to the errors in measurement of the H2 gas, like the solubility of H2 gas in water and the trapped H2 gas bubbles, and the fact that the molar volume is an average of two trials.
Trial 2’s result for molar volume ended up exceeding the 22. 43 L/mol accepted value, which could be due to error in reading the correct volume of the H2 gas in the tube or not allowing the pressure inside the tube to equalize completely to the atmospheric pressure. Overall, both trials are consistent in results that yield a molar volume of H2 gas relatively close to the accepted value of 22. 43 L/mol. To lessen error, the H? gas bubbles trapped in the liquid or on the sides of the tube should be gently shaken so that those bubbles can contribute to the measured volume of H? as. The eudiometer markings should be very precisely read to ensure that the measured volume is as close as possible to the actual volume. The pressure inside the tube after the reaction should be properly equalized with the atmospheric pressure to ensure the correct H? (g) volume measurement. The experiment should be performed multiple times, taking note of any temperature change or change in atmospheric pressure.
References: Chemistry by Zumdahl and Zumdahl, 6th ed. , N. p. : Houghton, 2003, n. d. 191. Print. Determining the Molar Volume of a Gas. N. p. : Houghton, n. . 87-92. Print. Vapor Pressure of Water at Different Temperatures. N. p. : Houghton, n. d. 88. Print. Appendix: Table of Contents: I. Pre-Lab Background Info and Questions A. Eva Leung B. Adrianne Pan C. Jane Kwong D. Jessica Kai II. Procedure Flow Charts A. Eva Leung B. Adrianne Pan C. Jane Kwong D. Jessica Kai III. Experimental Data/Observation Tables IV. Post-Lab Calculations and Questions (on same page as III) A. Eva Leung B. Adrianne Pan C. Jane Kwong D. Jessica Kai ————————————————- ————————————————-
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