Generating hydrogen gas lab answers

Experiment 5: GAS EVOLUTION REACTION Chemistry 121

I. Background / Theory

Many metals react with acids to produce hydrogen gas. These reactions are REDOX, gas evolution reactions. The metal atoms lose electrons to become metal ions. The hydrogen ions (from the acid) gain electrons to become hydrogen atoms, and two hydrogen atoms combine to form the hydrogen molecule, H2, which is a gas at room temperature. Under standard temperature and pressure, STP, all gases act like an ideal gas, thus following the ideal gas law with respect to the relationship between moles of the gas, pressure of the gas, temperature of the gas, and volume of the gas.

Based on the relationships of these different factors, the molar volume, or the volume of one mole of a gas, for an ideal gas at STP is 22.4 L per one mole of gas, or 22.4 L/mol. From experimentation, scientists found that the molar volume of many different gases at STP is very close to 22.4 L/mol, regardless of the type of gas involved.

In this experiment, magnesium metal is reacted with aqueous hydrochloric acid (HCl), producing hydrogen gas as one of its products. This experiment seeks to determine whether your experimental data is consistent with ideal gas behavior by calculating a molar volume of the hydrogen gas produced at STP and comparing this to the ideal molar volume of 22.4 L/mol.

In order to determine the molar volume of the hydrogen gas produced at STP (L/mol) from this reaction, the volume of the gas at STP (L) must be divided by the moles of the gas at STP (mol). Both of these values need to be determined. Unfortunately, this information can not be determined directly, since the experiment is not performed at STP.

However, the moles of gas at STP can be determined indirectly by recognizing the 1:1 stoichiometric relationship between the magnesium metal, the limiting reactant, and the hydrogen gas produced, assuming that all of the magnesium reacts.

Mg (s) + 2H+ (aq) -------------> Mg2+ (aq) + H2 (g)

Thus, by converting the mass of the magnesium metal into moles of magnesium and then into moles of hydrogen gas produced, the theoretical moles of gas at STP are calculated. The mass of the magnesium ribbon, however, can also not be directly determined, since the mass of the ribbon is too small to be measured on the scales provided in the laboratory. Instead, the magnesium ribbon mass is determined from the length of the magnesium ribbon itself and a pre-determined conversion factor, the linear density, that is provided. Since this process assumes that all of the magnesium ribbon gets used during the reaction, the ideal gas law can be used directly, solving for the actual moles of hydrogen gas produced. This value is then compared to the theoretical moles of hydrogen gas to validate that the magnesium did indeed get used up during the reaction.

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PV = nRT

Experiment 5: GAS EVOLUTION REACTION Chemistry 121

The determination of the volume of the gas at STP involves a more complicated calculation process. Since the experimental conditions are not at STP, the combined gas law is used to determine the volume of gas at STP.

This is done by assigning condition one as the experimental values and condition two as the STP values.

The volume of hydrogen gas at STP is then determined by rearranging the equation and solving for VSTP. The temperature and pressure at STP are standards and therefore are already known, 273.15 K and 1.00 atm, respectively.

Since the reaction is arranged so that the gas is collected over water in a sealed, graduated tube, the volume, VExp, can be read directly from the tube, as is done for a buret or a pipet, once the pressure is accounted for as described below.

Since the magnesium is placed in the tube so that the gas will bubble straight up displacing the water while filling the tube, the pressure of the gas above the water is actually a mixture of two gases, the hydrogen gas of interest and water vapor. As the hydrogen gas moves through the liquid water, some of the water molecules at the liquid-gas interface will acquire enough energy to escape into the gas phase. Although there is not a lot of water vapor present in the gas phase, the amount is large enough to affect the pressure and must therefore be corrected.

Dalton’s law of partial pressures is used to correct for this. By making the internal gas pressure (pressure inside the tube) equal to the external gas pressure (pressure of the atmospheric gases around the tube), the experimental pressure of the hydrogen gas can be determined. The inside and outside pressures are adjusted to be equal by raising or lowering the tube in water inside a larger graduated cylinder located in the sink so that the inside water level is equal to the outside level.

The external pressure is atmospheric pressure, or the room air pressure, which is read on the barometer. The internal pressure is the sum of the partial pressures of the experimental hydrogen gas and the partial pressure of the water vapor.

The vapor pressure of water can be read from a table posted on the wall of the lab, listing the vapor pressure of water at various temperatures. The pressure of the experimental hydrogen gas is then calculated by rearranging the above equation.

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P1V1 T1

= P2V2 T2

PExternal = PInternal

Patmospheric = PH2 + PH2O

PExpVExp TExp

= PSTPVSTP TSTP

PH2 = Patmospheric − PH2O

Experiment 5: GAS EVOLUTION REACTION Chemistry 121

The experimental temperature of the hydrogen gas is estimated by taking the average of the water temperature before the reaction, the water temperature after the reaction, and the room temperature.

With all of these calculations done, the molar volume of the hydrogen gas at STP can now be calculated. This calculated value is then compared to the theoretical value of 22.4 L/mol and a percent error can now be calculated as well.

II. Procedure:

Obtain a ribbon of magnesium metal and measure its length in centimeters (cm) using a ruler. The mass of the magnesium per length is written on the container of magnesium ribbons (g/m). Do not forget to record this conversion factor.

Ball up the ribbon and wrap it with copper wire to form a tight cage around the magnesium. Leave a small "tail" of copper trailing from this cage.

Set up a ring stand with a clamp or a buret stand to hold the 100 mL gas tube. Obtain the gas tube and a one-hole stopper. Fill a 400 mL beaker with distilled water. Place a thermometer in this water and record the temperature. Add about 10 mL of 6 M hydrochloric acid into the gas tube drop-wise, using the gas tube itself to measure the 10 mL. Slowly fill the tube to the brim with distilled water in such a manner as to disturb the acid layer as little as possible. The 400 mL beaker should be about half filled with water at this point.

Holding the copper cage by the tail, insert the cage into the tube filled with water and acid. Hook the tail over the edge of the tube, holding the case in place with you hand outside of the tube while you insert the one-hole rubber stopper. The tube and the hole in the stopper should be completely filled with liquid water.

Cover the hole in the stopper with your finger and invert the tube below the water level into the 400 mL beaker. Mount the tube vertically with the base just off the bottom of the beaker.

The more dense acid will fall down through the tube to the cage with the magnesium and react with it. Bubbles of hydrogen gas will rise and fill the tube, forcing the liquid water out of the tube through the hole in the stopper in a process known as displacement.

After the reaction stops, wait 5 minutes to allow the temperature in the tube to adjust back to room temperature. Then read the beaker water temperature. (The reaction is exothermic.) Cover the hole in the stopper with your finger and transfer the tube to the large cylinder by the sink, keeping the gas tube in the same vertical position (do not turn it over). Lower or raise the tube so the water levels in the tube and in the cylinder are the same. Read the gas volume to 0.1 mL.

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%Error = MolarVolumeTheoretical −MolarVolumeExperimental

MolarVolumeTheoretical ×100%

Experiment 5: GAS EVOLUTION REACTION Chemistry 121

After collecting the volume reading, the tube may be turned over and discarded properly.

Read the atmospheric pressure from the barometer. The units on the barometer are in cmHg.

Estimate the gas temperature by averaging the room and beaker temperatures (before and after reaction). From the table posted on the wall in the lab, look up the vapor pressure of water at this estimated temperature of the gas. Note that the units here are in mmHg. Do not forget to record a citation for this table. The citation is written on the table itself.

Pour the acid solution from inside the tube into the appropriate waste container in the chemical safety hood. Rinse the tube with distilled water into the same container and then return it to the cart to air dry. Unwind the copper wire, rinse it with distilled water into the waste container in the chemical safety hood, and then place it in the used copper wire bin on the back counter. Place all other equipment that came from a Chem 121 drawer back in the drawer. Wipe the counter clean. Wash your hands.

III. Qualitative and Quantitative Data:

1. Mass of Mg per meter (conversion factor in g/m). 2. Length of Mg ribbon in cm (to nearest 0.1 cm).

3. Room temperature (°C). 4. Beaker water temperature (°C): a. Before reaction.

b. After reaction.

5. Volume of gas collected (mL) . 6. Barometric pressure (cmHg). 7. Vapor pressure of water at your average temperature (mmHg; include reference).

Do not forget to include observations (1) before the reaction, (2) during the reaction, and (3) after the reaction.

IV. Calculations:

1. Mass of Magnesium Ribbon Reacted 2. Moles of Magnesium Ribbon Reacted 3. Theoretical Moles of Hydrogen Gas Produced

4. Atmospheric Pressure in mmHg 5. Partial Pressure of Hydrogen Gas in mmHg (Use Dalton’s Low of Partial Pressures) 6. Partial Pressure of Hydrogen Gas in atmospheres

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Experiment 5: GAS EVOLUTION REACTION Chemistry 121

7. Estimated Temperature of the Hydrogen Gas in ºC (Average temperature of the room temperature, the water temperature before, and the water temperature after).

8. Estimated Temperature of the Hydrogen Gas in K (Average temperature of the room temperature, the water temperature before, and the water temperature after).

9. Volume of Hydrogen Gas in L

10. Actual Moles of Hydrogen Gas Produced (Use the ideal gas law. This value should be similar to your answer in #3, assuming that all of the magnesium ribbon reacted.)

11. Ratio of Moles of Mg to Actual Moles of Hydrogen Gas Produced.

12. Volume that the Sample of Hydrogen Gas would Occupy at STP (Use the combined gas law). 13. Molar Volume of Hydrogen Gas at STP (Use the volume from #11 & the moles of hydrogen gas from

#3) 14. Using the theoretical value of 22.4 L/mole, calculate your % error.

V. Discussion Questions:

Consider the ideal gas law equation used to calculate the moles of H2 (#10 above): PV = nRT n = PV/RT

Referring back to this equation, explain what effect (increase, decrease, no change) the following procedural errors have on the value of n (moles of hydrogen gas):

a. The Mg did not react completely because a small piece floated out of the tube. b. The pressure of the hydrogen gas was recorded without subtracting the vapor

pressure of water. c. Some of the water in the tube was allowed to leak out through the hole in the stopper before

measuring the volume. d. The pressure was recorded based on the room temperature only, rather than averaging the three

different temperatures.

VI. Discussion of the Results

Include a general discussion on possible sources of error and how they may have influenced the data and/or final results. See the sample lab report and lab expectations for more help.

VII. Conclusion

Write a brief conclusion to wrap up the experiment. Did you successfully accomplish the objective? Why or why not? See the sample lab report and lab expectations for more help.

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