XHCO3(s) + HCl(aq) XCl(aq) + H2O(l) + CO2(g)

Transcription

1 Determining the Molar Mass of an Unknown Carbonate Using the Ideal Gas Law In this lab you will determine the molar mass of an unknown carbonate by using the ideal gas law to determine the number of moles of carbon dioxide produced when the carbonate reacts with hydrochloric acid. Knowing the moles of carbon dioxide will allow you to calculate the moles of the unknown. Combined with the mass of the unknown this will allow you to calculate the molar mass of the unknown carbonate. Stockroom You will need a quart size canning jar and a pint size canning jar with the appropriate glass tubing installed in the lids, pieces of rubber tubing, a pinch clamp, a digital thermometer, a 5mL burette, and a cuvette. The class will need several 500 ml graduated cylinders and a barometer. Equipment You will need your 600 ml beaker and your 150 ml beaker. Chemicals You will need about 5 grams of an unknown carbonate and about 100 ml of standardized 3M HCl. Introduction In this experiment you will react your unknown carbonate with hydrochloric acid. The general reaction is XCO3(s) + HCl(aq) XCl(aq) + HO(l) + CO(g) or XCO3(s) + HCl(aq) XCl(aq) + HO(l) + CO(g) or XHCO3(s) + HCl(aq) XCl(aq) + HO(l) + CO(g) depending on the cation. There are no unknowns with a +3 cation. The important thing here is that there is always a 1:1 mole ratio between the carbonate and carbon dioxide. This means that knowing the moles of carbon dioxide produced gives you the moles of your unknown. Because you will weigh your unknown, you will know both the mass and the number of mole of your unknown in that mass. This means you can calculate the molar mass of your unknown. You will measure the pressure, volume, and temperature of the carbon dioxide gas evolved. Using the ideal gas law this will give you the moles of carbon dioxide gas above the solution. Because some carbon Page 1 of 9

2 dioxide is dissolved in the solution we will take this into account also. Because you will be collecting the carbon dioxide gas over water, there will also be gaseous water molecules mixed in with the carbon dioxide gas molecules. This means the pressure we measure, the atmospheric pressure, will come from both carbon dioxide and water. You need to know the pressure of the carbon dioxide by itself. this use Dalton's law of partial pressures. To do You will read the atmospheric pressure from the barometer in the hallway. It looks like this: You will use the inner scale, which reads inches of mercury (inhg). The smallest marks are 0.0 inhg apart. The small numbers are 0.1 inhg. In this picture the atmospheric pressure reads as inhg. You will need to look up the vapor pressure of water at the temperature you measure for your carbon dioxide. Do that using the following table. Page of 9

3 Vapor Pressure of Water as a Function of Temperature From CRC Handbook of Chemistry and Physics 65th Edition T(oC) P water(torr) T(oC) Pwater(torr) T(oC) Pwater(torr) Page 3 of 9

4 Because CO is somewhat soluble in water, you will use the following equation (Equation 1) to calculate the mole fraction of CO dissolved in your solution (the Erlenmeyer flask) at the temperature of your reaction. You will know the partial pressure of CO above your solution because you used Dalton's law of partial pressures to calculate it. The following equation lets you calculate the mole fraction of carbon dioxide dissolved in your solution at the temperature of your solution (in oc) and the partial pressure of carbon dioxide above your solution (in kpa). X CO =( x 10 7 T x 10 5 T x 10 3 T ) P CO 5000 Equation 1 Here T is the temperature of your solution in oc and PCO is the pressure of CO in kpa in your flask. PCO =P atm PH O Where Patm is the atmospheric pressure that you read from the barometer in the lab and PH O is the vapor pressure of water at the temperature of your flask that you read from the above table. Be careful of units here. The atmospheric pressure you read from the barometer was in inhg, the table of water vapor pressures is in Torr, and you need kpa for Equation 1. To find the volume of carbon dioxide produced you will use the following glassware, set up as in the picture below. Page 4 of 9

5 Glassware for Finding the Volume of Gas Produced in a Chemical Reaction The chemical reaction occurs in the small jar on the left. The gas produced travels through the rubber tubing into the larger jar in the middle. The gas exerts a pressure on the water in the middle flask, causing the water to travel through the rubber tubing on the right. The water is collected in the 600 ml beaker on the right. The volume of water pushed over is the same as the volume of gas that is evolved in the chemical reaction, minus the gas that stays dissolved in the solution. You will measure the temperature of the water in the small jar (where the reaction took place) as soon as the reaction is complete. We can assume that the temperature of the gas is the same as the temperature of the solution. Page 5 of 9

6 Procedure 1.) Get your unknown from the stockroom. Record your unknown number in your data table and in your conclusion!.) Fill your large jar with tap water and put the lid on Connect the rubber tubing from the small jar to the glass large jar that does not down very far inside of the jar. get the inside of the tubing wet. DO NOT PUSH THE RUBBER FAR ONTO THE GLASS TUBE!!! tightly. tube in the It helps to TUBING TOO 3.) Insert a short piece of glass tubing into one end of a piece of rubber tubing. Put the other end of that rubber tubing on the glass tube in the large jar that extends almost to the bottom of the jar. DO NOT PUSH THE RUBBER TUBING TOO FAR ONTO THE GLASS TUBE!!! Place the end of rubber tubing that has the small piece of glass tubing into your 600mL beaker. 4.) Get a clean, dry cuvette and place it on the balance. Tare the balance. Add about 1.5 grams of your unknown to the cuvette and place it back on the balance. Record the mass of your unknown in your data table. 5.) Carefully place the cuvette with your unknown in it into the small jar making sure to not spill it. 6.) You might need to get help from another student for this step. One person push air through the rubber tubing from the glass tubing in the small jar. When water is coming out from the rubber tubing in the 600 ml beaker the other person clamps off the rubber tubing as close to the end as they can using a pinch clamp. 7.) Pour the water that was pushed into the 600 ml beaker back into the large jar. 8.) Place your 150 ml beaker on the balance and tare it. Using a burette add about 5 ml of standardized hydrochloric acid to your tared 150 ml beaker. Place the 150 ml back on the balance and record the mass of the hydrochloric acid in your data table. Record the volume of hydrochloric acid added to places past the decimal in your data table. Record the molarity of the hydrochloric acid in your data table (read it off of the label on the bottle). Page 6 of 9

7 9.) Pour the hydrochloric acid into the small jar, making sure it does not touch your unknown yet. 11.) Place the lid back onto the small jar. both jars are tightened securely. Make sure the lids on 10.) Remove the pinch clamp from the rubber tubing in the 600 ml beaker. 11.) Tilt the small jar so that your unknown carbonate spills into the hydrochloric solution. 1.) Continue to shake the small jar to make sure all of your unknown reacts. 13.) When the reaction is complete place the pinch clamp back on the rubber tubing in the 600 ml beaker. 14.) Immediately measure the temperature of the solution in the small jar using a digital thermometer. Record this in your data table. 15.) Read the atmospheric pressure in inhg in the hallway. this in your data table. Record 16.) Pour the water that collected in your 600 ml beaker into a 500 ml graduated cylinder. Record the volume of water in your data table. 17.) Repeat steps.) 16.) more times for a total of 3 trials. 18.) Rinse the small jar and lid three times with D.I. water. Page 7 of 9

8 Calculations Calculate the atmospheric pressure in torr. Calculate the partial pressure of CO above your solution by subtracting the vapor pressure of water from the atmospheric pressure. Calculate the partial pressure of CO above your solution in kpa. To calculate the moles of CO dissolved in your solution: Find the mole fraction of carbon dioxide ( X (CO ) ) in your solution nco from Equation 1 above. ( X CO )= where nco is moles of carbon n Total dioxide and ntotal is the total moles in the solution. That is, ntotal =nco + nh O +n HCl. However, some of the HCl reacted. If we take this into account we get ntotal =nh O +n HCl (initial) nco (above solution) nco (solution). However, we can neglect everything except nh O and nhcl(initial). This gives us: nco ( X CO )= Equation n H O +nhcl initial To find nhcl multiple the volume of HCl added, in liters, times the molarity of the standardized ~3M HCl you used. To find nh O find the mass of HCl added by multiplying nhcl by the molar mass of HCl and subtract this from the mass of HCl solution you weighed out. This is the mass of HO. Divide this by the molar mass of HO to get nh O. Use the ideal gas law to calculate the moles of CO collected above the solution and add to this the moles CO you calculated that was dissolved in your solution. Because the mole to mole ration is 1:1 for each carbonate:co we can say that moles of your unknown = moles of CO you just calculated. Calculate the molar mass of your unknown for each trial by dividing the mass you weighed out of your unknown by the moles of unknown calculated above. Page 8 of 9

9 Calculate the average molar mass of your unknown. Conclusion 1.) Your unknown number.) Your average molar mass of your unknown carbonate. 3.) Discuss thoroughly at least two possible sources of experimental error. You must include an analysis for each potential source of error that shows explicitly what effect that error would have on your final answer. Page 9 of 9