Chemistry Honors

Ch. 11 Packet ChemIsTry Honors

-Gas Stoichiometry -Avogadro’s Law -Diffusion and Effusion

Name:______Period:______19.1 Avogadros’ principle 1. What is Avogadro’s principle?

2. Based on Avogadro’s principle, if the volumes of two gases under similar conditions are equal, how are their number of moles related?

3. What is molar volume?

4. State what each of the five symbols in the ideal gas equation stands for

5. What temperature scale is used to solve problems involving the ideal gas equation?

6. Using the ideal gas equation, complete the table by solving for the unknown variable. Show your work below.

P (kPa) V (dm3) n T (K) #1 55.6 85.0 6.50 #2 210.0 3.20 293 #3 113 8.5 311 #4 95.1 5.50 277

7. In the modified form of the ideal gas equation PV = mRT/M, what do the symbols m and M stand for?

8. How can this modified equation also be used to solve for the density of gas?

9. Complete the table by solving for M. Show your work below

Atm P (kPa) V (dm3) m (g) T (K) M (g/mol) 0.973 98.6 6.27 0.832 300.0 1.11 112.5 10.1 2.36 298 0.875 88.7 23.5 0.55 292.5 0.900 91.2 2.75 0.124 0.312

Write T for true or F for false. If a statement is false, replace the underlined word or phrase with one that will make the statement true, and write your correction on the blank provided

10. ______At a given temperature_, the average kinetic energy of all gas molecules is the same. 11. ____,One __gram__of oxygen will occupy 22.4 dm3 at STP.

12. ___, The Celsius temperature scale is used for problems involving the ideal gas equation.

13. _____ Density is equal to mass divided by volume.

14. ____ The idea gas equation combines Boyle’s law and Charles’s law.

Gas Stoichiometry EXAMPLE: What volume of oxygen gas, O2, at STP can be produced when 17.5 grams of potassium chlorate is decomposed by heating? Solving process: In Section 9.2, we learned a method of solving mass-mass problems called stoichiometry. It is usually difficult to measure the mass of a gas. An easier method is to measure the volume under existing conditions and convert the volume to standard conditions. We have learned that 1 mole of any gas at STP occupies 22.4 dm³, the molar volume. We shall use this knowledge to determine the volume of gas in a reaction by using the balanced equation for the reaction. a. Write a balanced equation. 2KClO3  2KCl + 3O2 b. Express the mass (17.5g) of potassium chlorate, KClO3, in moles. 17.5 g KClO3 1 mol KClO3 123 g KClO3 c. Determine the mole ratio from the balanced chemical equation. Note that 2 moles of KClO3 yield 3 moles O2. 17.5 g KClO3 1 mol KClO3 3 mole O2 123 g KClO3 2 mol KClO3 d. Express the volume of oxygen gas, O2, in terms of dm³ by using the relationship that one mole of any gas has a volume of 22.4 dm³ at STP. 17.5 g KClO3 1 mol KClO3 3 mole O2 22.4 dm³ 123 g KClO3 2 mol KClO3 1 mol = 4.78 dm³ O2

7. How many dm³ of hydrogen gas at STP will be produced from 16.7 g of magnesium reacting with an excess amount of hydrochloric acid, HCL?

8. How many dm³ of carbon dioxide gas will be produced when 75.0 g of calcium carbonate, CaCO3, are decomposed into calcium oxide, CaO? (at STP)

9. Chlorine gas will react with 3.45 dm³ of hydrogen gas to yield what mass of hydrogen chloride gas, HCl, at STP? Gases W/S

Important: At the same temperature and pressure, equal volumes of gases have the same number of molecules. Therefore at STP, 22.4 L of any gas has 6.022 x 10²³ molecules (or atoms). This is obviously 1 mole.

All gases typically follow gas laws (variations between pressure, temperature and volume). Mathematically, these laws can be combined to from the ideal gas equation.

PV = nRT

There are two variations on the ideal gas equation that allow us to calculate density and molecular weight. Can you write those two equations below?

Practice problems

1) A NASA scientist proposed that barium peroxide, BaO2 (s) be used on a space capsule in order to supple emergency oxygen. It decomposes on heating according to the equation

2 BaO2 (s)  2 BaO (s) + O2 (g)

(a) What mass of BaO2 (s) would be needed to supply enough oxygen to fill a 10000L space capsule to a pressure of 0.20 atm at 25 ºC?

(b) How long would this oxygen last if the crew operating at 20 ºC used 1.00L min-¹ in respiration?

2) At STP the mass 1.00 L of a gas is 1.89g. What mass of gas will occupy 1.00 L at 200ºC and a pressure pf 1.25 atm?

3) Determine the amount of carbon dioxide, in liters at STP, that can be obtained by heating 1.00 kg of calcium carbonate, CaCO3 (s), which decomposes according to the equation

CaCO3 (s)  CaO (s) + O2 (g)

Chapter Review Problems

1. One mole of He has a mass of 4.0026 g and 1.000 dm3 of He (at STP) has a mass of 0.1787 g. Calculate the molar volume of helium.

2. A sample of gas has a mass of 1.248 g and occupies 300.0 cm3 at STP. What is the molecular mass of this gas?

3. From the volume, temperature, and pressure data given, calculate the ideal gas equation.

3 a) 2000.0 cm NH3 at 10.0 ºC and 105.0 kPa

3- b) 5.00 dm SO2 at 21.0 ºC and 100.0 kPa

4. Calculate the volume each gas will occupy under the conditions using the ideal gas equation.

a) 5.00 mol CH4 at 27.0 ºC and 97.2 kPa

b) 200.0 g NH3 at 12.0 ºC and 100.0 kPa

5. What is the molecular mass of a gas if 5.75 g of the gas occupy a volume of 3,50 dm3? The pressure was recorded as 9.525 x 104 Pa and the temperature s 52 ºC.

6. How many cubic centimeters of hydrogen at STP is produced by the reaction of 0.750 g of sodium metal with excess water?

2 Na (s) + 2 H2O (l) -> 2NaOH(aq) + H2 (g)

7. What mass of magnesium will react with excess hydrochloric acid to produce 2 3 5.00 x 10 cm of H2 at STP?

Mg (s) + 2HCl (aq) -> MgCl2 (aq) + H2 (g)

8. When lead(II) sulfide is burned in air, lead(II) oxide and sulfur dioxide are produced. If 20.0 dm3 of sulfur dioxide were produced, how many cubic decimeters of oxygen gas were required to react with the lead(II) sulfide?

2 PbS(s) + 3 O2 (g) -> 2 PbO (s) + 2 SO2 (g)

9. In a reaction involving carbon monoxide and iron(III) oxide, the products are iron metal and carbon dioxide. If 84.75 dm3 of carbon dioxide are produced, how many dm3 of carbon monoxide are required?

10. Calcium carbide (CaC2) reacts with water to produce calcium hydroxide and acetylene (C2H2). What volume of the gas at STP could be produced from the reaction of 50.0 g of CaC2 and 50.0 g of water? 3 3 11. Hydrogen burns to give water. If 200.0 cm of H2 reacts with 150.0 cm of O2, what volume of water vapor is produced? How many cubic centimeters of gas remain unreacted and what gas remains? Assume that all volumes are measured at any given temperature above the normal boiling point of water.

12. How many grams of sodium hydrogen carbonate, NaHCO3, must be heated to produce 2.50 dm3 of carbon dioxide measured at 22.5 ºC and 97.5 kPa? The other products are sodium carbonate and water.

13. If 3.20 g of aluminum react with excess hydrochloric acid, how many cm3 of hydrogen collected over water at 20.0 ºC and 99.5 kPa are produced The Reaction of Magnesium with Hydrochloric Acid

In this experiment you will determine the volume of hydrogen gas which is produced when a sample of magnesium reacts with 6 M hydrochloric acid. This is a solution of hydrogen chloride gas in water. The volume of the hydrogen gas produced will be measured at room temperature and pressure-conditions that matter when determining the volume of a gas. 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? Wear safety glasses and an apron.

PROCEDURE

a) Obtain a piece of magnesium ribbon approximately 2.5 cm long. Obtain the mass and cut it into four pieces. b) Look at Figure 12-1 to see how the magnesium ribbon is encased in a cage of fine copper wire. Notice that the cage has no large openings through which small pieces of magnesium ribbon could escape. c) Set up a ring stand and utility clamp in position to hold a 50 mL gas-measuring tube as shown in Figure 12-3. Fill a 400 mL beaker about two thirds full of tap water. Place it near the ring stand. d) Incline the gas-measuring tube slightly from an upright position and pour in about 5 mL of 6 M hydrochloric acid. e) With the tube 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 test tube so that the liquid in the top of the tube will contain very little acid. Try to avoid stirring up the acid layer in the bottom of the tube. Bubbles clinging to the side of the tube can be dislodged by tapping the tube gently. f) Holding the copper coil by the handle, insert the cage about 5 cm down into the Tube. See Figure 12-2 hook the copper wire over the edge of the tube and hold it there by inserting the rubber stopper. The tube should be completely filled so that the stopper displaces a little water when put in place. g) Cover the hole in the stopper with your finger and invert the tube in the container of water, as shown in Figure 123. Clamp it in place. The acid, being more dense than water, will fall down through it and eventually react with the metal. h) After the reaction stops, wait for about 5 minutes to allow the tube to come to room temperature. Dislodge any bubbles clinging to the side of the tube. i) Cover the hole in the stopper with your finger and transfer the tube to a large cylinder or battery jar which is almost filled with water at room temperature. See Figure 12-4. Lower or raise the tube until the level of liquid inside the tube is the same as the level outside the tube. This permits you to measure the volume of gases in the tube (hydrogen and water vapor) at room pressure. Read the gas volume. Your eye should be at the same level as the bottom of the meniscus (the lens-shaped surface taken by the water in tube). See Figure 12-5. Record the volume of the gas as precisely as you can. j) Remove the gas-measuring tube from the water and pour the acid solution it contains down the sink. Rinse the tube with tap water.

k) k) Record room and water temperatures. Your teacher will give you the room pressure or will assist you in reading the barometer.

PROCESSING THE DATA 1) Determine the mass of the magnesium you used knowing the grams per meter and the length of your piece of ribbon. 2) Determine the number of moles of magnesium used. 3) Determine the partial pressure 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 i. See the hypothetical case illustrated in Figure 12-6A. Mathematically this can be expressed PH2 + PH2O = Proom The pressure of the room may be determined by reading the barometer. The pressure of the water vapor, PH2O, can be obtained from the table given above. The values in the table were obtained by measuring the pressure of water vapor above liquid water at various temperatures. The partial pressure of the hydrogen can be calculated as follows: PH2 = Proom - PH2O

Figure 12-6A

4) Determine the volume the hydrogen gas would have at 1 atmosphere pressure (760 mm Hg). You have learned that for a given temperature the product of the pressure and volume of a gas is a constant. To calculate the new volume, Vnew, at 760 mm Hg pressure, the following mathematical relation can be stated: PH2 x Vmeasured = 760 x Vnew Or, solving for Vnew Vnew = Vmeasured x PH2/760

Fig. 12-5> Reduce errors by reading the bottom of the meniscus with your eye at the proper level

5) Calculate the volume of dry hydrogen which would be produced by 1 mole of magnesium at room temperature and 1 atmosphere pressure. Use your data. 6) Write a balanced chemical equation for the reaction Mg + HCl. Given that 1 mole of Mg produces 1 mole of hydrogen, H2, what is the volume of 1 mole of hydrogen at room temperature and 1 atmosphere pressure? (Use your data.) 7) If 1 mole of hydrogen weighs 2.0 g, what is the mass of a liter (the density) of hydrogen at room temperature and 1 atmosphere pressure? ADDITIONAL INVESTIGATION Consult your teacher before proceeding. Determine the volume of hydrogen gas produced when a mole of another metal reacts with an acid. YOUR DATA TABLE SHOULD INCLUDE THE FOLLOWING: Mass of magnesium ribbon in grams per meter (from teacher) Length of magnesium ribbon used Volume of hydrogen and water vapor collected Temperature of the water Temperature of the room Barometer reading (room pressure) Vapor pressure of water at the above temperature (see the following table)