Determining the Molar Mass of an Unknown Carbonate Using the Ideal Gas Law

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 1 Florence flask with 2 hole stopper and glass tubing inserted, a 1 hole rubber stopper for your 250 ml Erlenmeyer flask with glass tubing inserted, 2 pieces of rubber tubing, a pinch clamp, a digital thermometer, and 3 cuvettes. The class will need several 500 ml graduated cylinders and a barometer. Equipment You will need your 600 ml beaker, your 150 ml beaker, and your 250 ml Erlenmeyer flask. Chemicals You will need about 6 grams of an unknown carbonate and about 35 ml of 6M HCl (with a 10 ml graduated cylinder by the bottle). Introduction In this experiment you will react your unknown carbonate with hydrochloric acid. The general reaction is X 2 CO3 (s)+2 HCl( aq ) 2 XCl( aq )+H 2 O (l )+CO 2 ( g ) or XCO3 (s)+2 HCl( aq ) XCl2 ( aq )+ H 2 O( l )+CO2 (g ) depending on the charge on the cation. +3 cation. There are no unknowns with a 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. NAME: 1 of 14

2 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. Molar mass of unknown = Mass of unknown Moles of unknown (Equation 1) To find the moles of carbon dioxide produced by your unknown you will use the ideal gas law. PV =nrt or PV n= RT (Equation 2) The variables in this equation and their units are: P=pressure in of the gas in atm V=volume of the gas in liters L atm R= K mol T=temperature of the gas in Kelvins You will measure the pressure, volume, and temperature of the carbon dioxide allowing you to calculate the moles of the carbon dioxide, which gives you the moles of your unknown. Because you will be collecting the carbon dioxide gas over water, there will also be gaseous water molecules mixed in with the carbon dioxide molecules. This means the pressure we measure, the atmospheric pressure, will come from both carbon dioxide and water. We need to know the pressure of the carbon dioxide by itself. To do this, we will subtract the pressure of the water from the total pressure (which is the same as the atmospheric pressure). Patm = P CO +PH O or PCO = Patm PH O (Equation 3) 2 You will read the atmospheric pressure from the barometer in the hallway. It looks like this: NAME: 2 of 14

3 You will use the inner scale, which reads inches of mercury (inhg). The smallest marks are 0.02 inhg apart. The small numbers are 0.1 inhg. In this picture the atmospheric pressure reads as inhg. The number you read from the barometer is the atmospheric pressure (Patm). To use equation 3 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. To find the vapor pressure of water at the temperature of your gas, find the temperature to the integer place in the left column and the tenths place in the top row. For example, if your temperature is 37.3 oc then the vapor pressure of water is 47.7 mmhg. The integer value (37) is the blue box, the tenths (0.3) is the yellow box, and the pressure (47.7 mmhg) is the green box. NAME: 3 of 14

4 Vapor Pressure of Water in mmhg from the Antoine Equation T (oc) To find the volume of carbon dioxide produced you will use the following glassware, set up as in the picture below. NAME: 4 of 14

5 Glassware for Finding the Volume of Gas Produced in a Chemical Reaction The chemical reaction occurs in the Erlenmeyer flask on the left. The gas produced travels through the rubber tubing into the Florence flask 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. You will measure the temperature of the water in the Erlenmeyer flask on the left as soon as the reaction is complete. We can assume that the temperature of the gas is the same as the temperature of the water. NAME: 5 of 14

6 Procedure 1.) Get your unknown from the stockroom. the data section at the top of page 8. Tape your unknown number in 2.) Fill your Florence flask with tap water and insert the two hole rubber stopper. Connect the rubber tubing to the glass tubing in the two hole rubber stopper. It helps to get the inside of the tubing wet. 3.) Connect the piece of rubber tubing that is connected to the short piece of glass tubing in the Florence flask to the one hole stopper for the Erlenmeyer flask. The piece of rubber tubing that is connected to the long piece of glass tubing in the Florence flask should be placed in your 600 ml 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 the data section (A1), (B1), (C1). 5.) Carefully place the cuvette with your unknown in it into the Erlenmeyer flask making sure to not spill it. 6.) You will need to get help from another student for this step. One person push air through the rubber tubing from the glass tubing in the one hole rubber stopper. 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 Florence flask. Take the Florence flask, rubber stoppers, and rubber tubing over to a sink and make sure the Florence is filled completely. Over the sink insert the two holw rubber stopper into the Florence flask. Bring this back to your lab bench. 8.) Obtain a little more than 10 ml of 6M hydrochloric acid. 9.) Pour the hydrochloric acid into the Erlenmeyer flask, making sure it does not touch your unknown yet. 10.) Place the one hole rubber stopper into the Erlenmeyer flask. Make sure that both rubber stoppers are in securely. You might need NAME: 6 of 14

7 to hold them to keep the rubber stoppers tight during the reaction. 11.) Tilt the Erlenmeyer flask so that your unknown carbonate spills into the hydrochloric solution. 12.) Immediately remove the pinch clamp from the rubber tubing in the 600 ml beaker. MAKE SURE TO HOLD BOTH RUBBER STOPPERS DOWN SO THAT THE GAS DOESN'T LEAK OUT! 13.) Continue to shake the Erlenmeyer flask to make sure all of your unknown reacts. 14.) When the reaction is complete place the pinch clamp back on the rubber tubing in the 600 ml beaker. 15.) Measure the temperature of the water in the Florence flask using a digital thermometer. Record this in the data section (A2), (B2), (C2). 16.) Read the atmospheric pressure in inhg in the hallway. this in the data section (A3), (B3), (C3). Record 17.) Pour the water that collected in your 600 ml beaker into a 500 ml graduated cylinder. Record the volume of water in the data section (A4), (B4), (C4). 18.) Repeat steps 2.) 17.) 2 more times. NAME: 7 of 14

8 Data and Analysis DATA Unknown Number Mass of unknown 1st trial (A1) Temperature 1st trial (A2) Atmospheric pressure 1st trial (A3) Volume of water collected 1st trial (A4) Mass of unknown 2nd trial (B1) Temperature 2nd trial (B2) Atmospheric pressure 2nd trial (B3) Volume of water collected 2nd trial (B4) Mass of unknown 3rd trial (C1) Temperature 3rd trial (C2) Atmospheric pressure 3rd trial (C3) Volume of water collected 3rd trial (C4) NAME: 8 of 14

9 Analysis Show all work including significant figures and units. Record your answer in the blank provided to the correct number of significant figures with the correct units. 1.) Convert the temperature from trial 1 to Kelvins. (A2) (A5) 2.) Convert the atmospheric pressure from trial one to mmhg. 760 mmhg (A3) (29.92 inhg) (A6) 3.) Find the vapor pressure of water at the temperature of your reaction from the table on page 4. Record it here. (A7) 4.) Find the vapor pressure the CO2 by subtracting the vapor pressure of water from the atmospheric pressure. (A6) (A7) (A8) 5.) Convert the pressure of your CO2 in trial 1 to atmospheres. 1 atm (A8) 760 mmhg (A9) NAME: 9 of 14

10 6.) Convert the volume of water collected in trial 1 (which is the same as the volume of CO2) to liters. 1L (A4) 1000 ml (A10) 7.) Using the ideal gas law (Equation 2) calculate the moles of CO2 produced in the first trial. This is equal to the moles of your unknown that you reacted. (A9)(A10) (R)(A5) (A11) 8.) Calculate the molar mass of your unknown for trial 1 using (Equation 1). (A1) (A11) (A12) NAME: 10 of 14

11 9.) Convert the temperature from trial 2 to Kelvins. (B2) (B5) 10.) Convert the atmospheric pressure from trial two to mmhg. 760 mmhg (B3) ( inhg) (B6) 11.) Find the vapor pressure of water at the temperature of your reaction from the table on page 4. Record it here. (B7) 12.) Find the vapor pressure the CO2 by subtracting the vapor pressure of water from the atmospheric pressure. (B6) (B7) (B8) 13.) Convert the pressure of your CO2 in trial 2 to atmospheres. 1 atm (B8) 760 mmhg (B9) NAME: 11 of 14

12 14.) Convert the volume of water collected in trial 2 (which is the same as the volume of CO2) to liters. 1L (B4) 1000 ml (B10) 15.) Using the ideal gas law (Equation 2) calculate the moles of CO2 produced in the second trial. This is equal to the moles of your unknown that you reacted. (B9)(B10) (R)(B5) (B11) 16.) Calculate the molar mass of your unknown for trial 2 using (Equation 1). (B1) (B11) (B12) NAME: 12 of 14

13 17.) Convert the temperature from trial 3 to Kelvins. (C2) (C5) 18.) Convert the atmospheric pressure from trial three to mmhg. 760 mmhg (C3) (29.92 inhg) (C6) 19.) Find the vapor pressure of water at the temperature of your reaction from the table on page 4. Record it here. (C7) 20.) Find the vapor pressure the CO2 by subtracting the vapor pressure of water from the atmospheric pressure. (C6) (C7) (C8) 21.) Convert the pressure of your CO2 in trial 3 to atmospheres. 1 atm (C8) 760 mmhg (C9) NAME: 13 of 14

14 22.) Convert the volume of water collected in trial 3 (which is the same as the volume of CO2) to liters. 1L (C4) 1000 ml (C10) 23.) Using the ideal gas law (Equation 2) calculate the moles of CO2 produced in the third trial. This is equal to the moles of your unknown that you reacted. (C9)(C10) (R)(C5) (C11) 24.) Calculate the molar mass of your unknown for trial 3 using (Equation 1). (C1) (C11) (C12) 25.) Calculate the average molar mass of your unknown. (A12) + (B12) + (C12) 3 AVERAGE MOLAR MASS OF UNKNOWN: TURN IN PAGES 8 14 ONLY! NAME: 14 of 14