Lab - Reaction of Magnesium with Hydrochloric Acid

The Reaction of Magnesium with Hydrochloric Acid Objective In this experiment you will determine the volume of hydrogen gas that is produced when a sample of magnesium reacts with 6M hydrochloric acid. Introduction The volume of hydrogen gas produced when a magnesium sample reacts with HCl will be measured at room temperature and pressure. The data you obtain will enable you to answer the question: How many liters of dry hydrogen gas at room temperature and 1 atmosphere can be produced per mole of magnesium metal used? Pre Lab Problems (answer on separate paper, showing all appropriate work) 1. Write the balanced net ionic equation for the reaction of magnesium metal with hydrochloric acid. 2. If you used 0.500 g of Mg and an excess of HCl, how much hydrogen gas should be produced at STP? 3. What is the water vapor pressure at the above conditions? (This will require some searching, either open literature or other logical spots – make sure you properly cite your source). Procedure 1. Obtain a piece of magnesium ribbon approximately 5 cm long. Carefully determine the mass of the ribbon and record it. 2. Look at Figure 1 (right), which illustrates how the magnesium ribbon is encased in a cage of fine copper wire. Notice the cage has no large openings through which small pieces of magnesium ribbon could escape. Create a cage around your ribbon, making sure that it will fit inside a buret when complete. 3. Set up a ring stand and buret clamp in position to hold a buret as shown in Figure 2. Fill a 400-mL beaker about two-thirds full of tap water. Place it near the ring stand. 4. Incline the buret slightly from an upright position and pour in ~ 10 mL of 6M hydrochloric acid. Make sure the buret stopcock is closed before adding the liquid! 5. With the buret in the same position, slowly fill it with tap water from a beaker or bottle. While pouring, rinse down any acid that may be on the side of the buret so that the liquid at the top of the buret contains very little acid. Avoid stirring up the acid layer in the bottom of the buret. Bubbles clinging to the side of the buret can be dislodged by gently tapping the buret. 6. Holding the copper cage by the handle, insert the cage about 3 cm down into the buret (see Figure 3). After insertion, the cage should end up about where the 0 to 2 or 3 mL marking is on the buret. Hook the copper wire over the edge of the buret and hold it there by inserting the rubber stopper. The buret should be completely filled so that the stopper displaces a little water when put in place. 7. Covering the hole in the stopper with your finger, invert the buret in a beaker of water, (Figure 2). Clamp it in place. Hydrochloric acid has a greater density than water and will “fall” down, eventually reacting with the metal. 8. After the reaction stops (when there is no longer any bubbling), wait for about 5 minutes to allow the buret to come to room temperature. Dislodge any bubbles clinging to the side of the buret. 9. Cover the hole in the stopper with your finger and transfer the buret to a large cylinder which is almost filled with water at room temperature (see Figure 4). Lower or raise the buret until the level of the liquid inside the buret is at the same level as the water outside the buret. This permits you to measure the volume of the gases (hydrogen and water vapor) in the buret at room pressure. If the inverted buret cannot match liquid level within the buret with the liquid level in the large cylinder, it will be necessary for you to re-run the reaction. 10. Read the gas volume. Your eye should be level with the bottom of the meniscus. Record the volume of the gas to the nearest 0.01 mL. You will have to subtract the reading from 50.00 mL and then ADD the volume between the 50 mL mark and the top of the stopcock on the buret. Do this by adding water to the buret and letting it drain down until it is exactly at the 50 mL mark. Then empty the water remaining into a 10 mL graduated cylinder, and record this amount. This is the volume that you were instructed to ADD (the volume between the 50 mL mark and the top of the stopcock). 11. Remove the buret from the water and dispose of the acid solution down the sink. 12. Record both the room temperature and the barometric pressure. 13. Repeat the experiment with another sample of magnesium to check your results. If the two experiments don’t yield fairly close results (± 5%), you will need to conduct a third run. Cleanup Clean your lab area, glassware, and buret before being signed out. You will be required to construct your own Data Table for this lab. This data table will constitute a sizable portion of your grade for this lab. It should be well-organized, neat, legible, and should include the following: Mass of magnesium ribbon in grams Volume of hydrogen and water vapor collected Temperature of the room Temperature of the large cylinder water bath Barometric pressure Vapor pressure of water at the above temperature (see Table 1) Table 1 Vapor pressure of water at various temperatures Temperature (°C) 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 Pressure (mm Hg) 12.8 13.6 14.5 15.5 16.5 17.5 18.6 19.8 21.0 22.4 23.7 25.2 26.7 28.3 30.0 31.8 Post Lab Problems (answer - with calculations - on separate paper) 1. Determine the number of moles of magnesium used. 2. Determine the partial pressures of the hydrogen gas. Since the hydrogen gas was collected over water, the gas in the tube consists of a mixture of hydrogen gas and water vapor. The total pressure caused by these two gases was made equal to room pressure in Step 10, so: PH2 + PH2O = Proom 3. Determine the volume the dry hydrogen gas would have at one atmosphere pressure (760 mm Hg). You now have the volume of dry hydrogen gas at standard pressure, but not standard temperature. 4. Now convert the volume to standard temperature (0 °C). This should give you the volume of your hydrogen gas at STP (which is small because you used a small number of grams of magnesium). 5. Calculate the volume of dry hydrogen you would expect to produce by using 1 mole of magnesium at one atmosphere pressure and room temperature. (HINT: This should be a simple dimensional analysis conversion from the volume of dry hydrogen gas you got by using less than a gram of magnesium (see Problem 3) to what you expect if you had used 1 mole of magnesium). 6. Convert your value of the volume of hydrogen expected from 1 mole of magnesium (see Problem 5) to STP, and make a comparison between your value and the standard value for 1 mole of a gas at STP. 7. Determine the density of your hydrogen at STP. Compare it with the known density of hydrogen gas at STP. Remember, density is mass divided by volume.