Unit 12: The Behavior of Gases. (Chapter 14)

Transcription

1 Unit 12: The Behavior of Gases (Chapter 14) 1

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3 Gas Pressure A gas pressure you are familiar with is that caused by a mixture of gases the air! Air is a mixture of several gases and the gaseous molecules are in constant motion. This constant motion causes them to collide with objects and exert a pressure. This pressure is known as atmospheric pressure. We cannot see or touch atmospheric pressure we can only feel it when the wind is blowing! As the particles of gas decrease, the collisions are not as frequent, and pressure decreases. For example, atmospheric pressure decreases as you climb a mountain because the density of the Earth s atmosphere decreases as elevation increases. Pressure is the measure of force per unit area. SI unit for pressure is the. Two other units are still commonly used: o millimeters of mercury o atmospheres *In the case of gases, it is important to be able to relate measured values to standards.* Standard pressure = Standard temperature and pressure ( ) = 1 atm = PRESSURE CONVERSIONS mm Hg torr kpa (kilopascals) psi (pounds per square inch) 3

4 Directions: Fill in the correct pressure conversions below. Then use the factor label method to solve the pressure conversion problems. 1 atm = mm Hg = torr = kpa = psi 1. Convert 0.75 atmospheres to torr. 2. Convert 450 kpa to atmospheres. What is this pressure in mmhg? 3. What pressure, in pounds per square inch, does a gas exert at 385 mm Hg? What is this pressure in kilopascals? 4. The pressure at the top of Mount Everest is 33.7 kpa. Is that pressure greater or less than 0.25 atm? 4

5 PRESSURE CONVERSIONS Many different units are available to use when taking measurements. When measuring length, you can use inches, feet, yards, or miles. The unit that is used usually best fits the situation. For instance, the distance from NASH to NAI would not be measured in inches, but in miles. If for some reason you needed this distance in inches, you could convert the miles to inches using the FACTOR LABEL METHOD. This is also true for pressure measurements. Different instruments, calibrated in different units, are used to measure the pressure of gases. Using the factor label method you can easily convert from one unit to another. 1 atm = 760 mm Hg = 760 torr = kpa = 14.7 psi (pounds per square inch) 1. On the average, the atmospheric pressure in Denver, Colorado, is atm. Express this pressure in (a) millimeters of mercury (mm Hg), and (b) kilopascals (kpa). 2. Because Pittsburgh is not at sea level, the atmospheric pressure is never 760 mm Hg, but averages around 730 mm Hg. Express this pressure in atmospheres (atm). 3. During a hurricane the atmospheric pressure can drop 100. mm Hg. Assuming that the normal atmospheric pressure in Miami is 1.00 atm, what could the pressure be in pounds per square inch (psi) during a hurricane? 4. The barometer reading that the weathermen use is expressed in inches of mercury (in Hg). What would a barometer reading of 100. kpa be in inches of Hg? (hint: 1 in = 2.54 cm) 5

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7 Pressure, Temperature, Volume and Mole Relationships W e live at the bottom of a sea of molecules. Molecular nitro-gen and oxygen make up 99% of the earth s atmosphere, but other gases are important as well. Carbon dioxide, for example, is needed by plants to produce sugars; and a layer of ozone gas, O 3, surrounds the planet and protects us from dangerous radiation from the sun. In many ways the gaseous phase of matter can be understood more easily than the liquid and solid phases. Under ordinary pressure conditions, gas molecules have little influence on one another. In contrast, the molecules of a liquid or solid are always close enough to have a strong influence on their neighbors. Four quantities are needed to define the state of a gas: (a) pressure of a gas, P; (b) temperature of the gas, T; (c) volume of the gas, V; and quantity of gas, n. In the 17th and 18th centuries, scientists experimented with gases and provided a basic understanding of their behavior. You are going to duplicate some of their work to discover the relationships among P, V, T, and n. BOYLE S LAW Robert Boyle ( ), the father of chemistry, was one of the first scientists to practice experimentation rather than just reason to explain scientific phenomena. His experiments with gases also initiated the practice of carefully describing one s work so that anyone could repeat and confirm it. His most famous work was his study of the relationship between the pressure and volume of a gas. In order to understand and to be able to describe this relationship, you are going to conduct a few experiments. ASSIGNMENT #1: Place a balloon inside a bell jar. Predict what will happen to the balloon in the vacuum. Turn on the vacuum pump to reduce the pressure inside the bell jar. Describe what happens to the balloon. Using the data from the demonstration performed above, describe what kind of relationship exists between the pressure and volume of a gas. (Remember direct and inverse?) Fun And Relevant Topics about Gases Write Boyle s Law by filling in the spaces below. At constant and, the and of a gas are related. CHARLES LAW Jacques Charles ( ) constructed the first hydrogen balloon in which he went up several times, reaching a height of over a mile and creating an aeronautical craze. He is best known for his study of the relationship between volume and temperature of a gas. From his experiments, he predicted absolute zero, the temperature at which the volume of a gas would be zero. Since he did not publish his results, it is sometimes called Gay-Lussac s Law after another balloon-ascensionist who duplicated Charles work. ASSIGNMENT #2: Obtain mL beakers and half-fill each with water. Blow up a balloon so that it just fits in the 600-mL beaker. Heat the water in one of the beakers to boiling; add ice to the second beaker to cool down the water. Predict what will happen to the balloons in each beaker. room temp hot cold Place the balloon in the room temperature beaker, then in the hot beaker, and finally in the cold beaker. Describe what occurred after each temperature change. room temp hot cold 7

8 Using the data from the demonstration performed above, describe what kind of relationship exists between the temperature and volume of a gas. Write Charles Law by filling in the spaces below. At constant and, the and of a gas are related. ASSIGNMENT #3: With Boyle s Law, any unit of pressure can be used. However, Charles found that the temperaturevolume relationship was dependent upon the temperature scale. He calculated absolute zero, which was later substantiated by Lord Kelvin in Write an equation to convert C to K. GAY-LUSSAC S LAW The French chemist Joseph Gay-Lussac ( ) was also an avid balloonist who ascended to a height of four miles in one of his flights; this record stood for many years. In one of his flights, he threw an old kitchen chair overboard to lighten the balloon and gain height. It landed near a peasant girl who was minding her sheep. After much discussion, the local citizens and priest decided that the incident was a miracle, but could not understand why God owned such shabby furniture. While studying gases, he found that in forming compounds, gases combine in proportions by volume that could be expressed in small whole number ratios. This ratio happens to be the same as the mole ratio for the reaction. He also studied how the pressure of a contained gas changed with temperature. ASSIGNMENT #4 Place 5 ml of water into a 1000-mL Erlenmeyer flask. Boil the water until it is almost all gone and the flask is filled with water vapor. Remove the flask from the ring stand. Place a peeled hard-boiled egg on top of the flask. Predict what will happen to the egg. As the flask cools, observe what happens to the egg. Write a scientific explanation for this phenomenon. Describe what kind of relationship exists between temperature and pressure. Using Boyle s and Charles Laws as models, write Gay-Lussac s Law. Using Gay-Lussac s Law, describe a method to get the egg out of the flask. AVOGADRO S LAW Gay-Lussac s work on combining volumes was ignored, particularly by John Dalton (remember him?), but Amadeo Avogadro ( ) expounded on Gay-Lussac s theory to state that all gases at the same temperature and pressure contain the same number of particles per unit volume. This, too, was refuted at the time, and Avogadro suffered the not-too-uncommon fate of neglect during his lifetime and success after death. Now, of course, Avogadro is famous, with his name applied to the number of particles in a mole 6.02 x ASSIGNMENT #5 Obtain a balloon and blow into it once. Observe its relative volume, temperature, pressure, and number of molecules. V = P = T = n = Of the four conditions T, P, V, and n which two radically change as you blow into a balloon? 8

9 Blow into it a second time and again record its relative volume, temperature, pressure, and number of moles. V = P = T = n = ASSIGNMENT #6 After looking at the pictures below, state the following: (a) the two variables and how they are changing, (b) the two constant conditions, and (c) the law. Repeat the procedure two more times until you see a pattern of V, T, P, and n develop. V = P = T = n = V = P = T = n = Describe what kind of relationship exists between the volume and number of particles of a gas. Using Boyle s and Charles Laws as models, write Avogadro s Law. TIME TO REVIEW You have just defined the relationships among the four conditions that describe a gas T, V, P, and n. In each one of these relationships, two of the conditions are variables (i.e., they change), while the other two conditions must remain constant. Picture 1 Volume increases remain constant Law Picture 2 Volume increases remain constant Law 9

10 What is the effect on the volume of one mole of gas when each of the following occurs? Briefly explain your answer in each case. a. The pressure is tripled. b. The Kelvin temp. is increased by a factor of 2.5. c. Two more moles of gas are added (at constant T and P). Picture 3 Volume increases remain constant Law What is the effect on the pressure of one mole of gas when each of the following occurs? Briefly explain your answer in each case. a. (The temperature is increased from 350 K to 700 K. b. The volume is decreased from 8 L to 2 L. c. Half the gas escapes through a stopcock. Picture 4 Volume increases remain constant Law 10

11 Review of Gas Laws BOYLE S LAW: Variables: Constants: LAW: At constant and, and of a gas are related. PRACTICE PROBLEM: The pressure on 2.50 L of anesthetic gas is changed from 1760 mm Hg to 760. mm Hg. What will the new volume be if the temperature remains constant? CHARLES LAW: Variables: Constants: LAW: At constant and, and of a gas are related. PRACTICE PROBLEM: If a balloon occupies 6.8 L at C, what will its volume be At C if the pressure does not change? 11

12 GAY-LUSSAC S LAW: Variables: Constants: LAW: At constant and, and of a gas are related. PRACTICE PROBLEM: The gas left in a used aerosol can is at a pressure of 1 atm at 27 0 C (room temperature). If this can is thrown onto a fire, what is the internal pressure of the gas when its temperature reaches C? AVOGADRO S LAW: Variables: Constants: LAW: At constant and, and of a gas are related. PRACTICE PROBLEM: A sample of 0.55 moles of a gas occupies 2.5 L at 755 mm Hg. If 0.45 moles of the gas are added to the balloon, what will be the new volume of the balloon? (assume constant pressure). 12

13 DIRECTIONS: Solve the problems. Label all of the given data and the unknown. 1. When asked to state Boyle's Law, a student replied, The volume of a gas is inversely related to the pressure. How is this statement incomplete? Give a correct statement of Boyle s Law. 2. A sample of ammonia gas occupies a volume of 450 ml at a pressure of 724 mm Hg. What volume will it occupy at standard pressure? 3. A sample of carbon dioxide gas occupies a volume a 3.50 liters at 125 kpa pressure. What pressure, in atm, would you exert if the volume was decreased to 2.00 liters? 4. A 2-L bag of Snyder s pretzels that was manufactured in Pittsburgh where the pressure is 740 mm Hg is taken to Pike s Peak where the pressure is only 84.6 kpa. What is the volume of the bag at this new altitude? 13

14 1. A 0.25-L air filled balloon at room temperature, 23 C, is placed in a beaker of liquid nitrogen at -196 C. What is the volume of the balloon while it is submerged in the N 2? What will the balloon do when it is removed from the liquid N 2? 2. A balloon is inflated with helium to a volume of 4.5 liters at room temperature, 25 C. If you take the balloon outside on a cold day (-10 C), what is the new volume of the balloon? 3. To what temperature would a mass of hydrogen gas have to be changed if 2.5 liters at 25 C is to be put into a 1.5 L container where the pressure remains constant? 4. A 2-L bag of Snyder s pretzels that was manufactured during the winter in a cold plant, 59 F, is taken to a summer picnic where the temperature is a scorching 95 F. What is the volume of the bag of pretzels? Do temperature changes have as great an effect on volume as pressure changes? (refer to your answer in #4 of Boyle s Law) 14

15 1. The pressure of a container of helium is 650.torr at 25 C. If the sealed container is cooled to 0 C, what will the pressure be? 2. The pressure in an automobile tire is 2.0 atm at 27 C. At the end of a journey on a hot sunny day, the pressure has risen to 2.2 atm. What is the temperature of the air in the tire? (Assume the volume has not changed.) 3. An iron tank of helium contains the gas at a pressure of 2000 psi at 25 C. The tank will hold pressures of about 7000 psi before exploding. If the tank is in a building that catches fire, will the tank explode before it melts? The melting point of iron is 1535 C. 1. One mole of any gas occupies a volume of 22.4-L. How many moles of gas do you need to fill a 1.0-L balloon? 15

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17 Review of Gas Laws BOYLE S CHARLES GAY-LUSSAC s Let s COMBINE the 3 Gas Laws P 1 V 1 = P 2 V 2 V 1 = V 2 P 1 = P 2 T 1 T 2 T 1 T 2 The Combined Gas Law 17

18 The Combined Gas Law 1. The volume of gas is 27.5 ml at 22 0 C and atm. What will the volume be at 15 0 C and atm? 2. A 700 ml gas sample at STP is compressed to a volume of 0.2 L, and the temperature is increased to C. What is the pressure of the gas in kilopascals? In torr? 18

19 Combined Gas Law DIRECTIONS: Answer the following questions in the space provided. Label all of the given data and the unknown. 1. Some octane is placed in the cylinder of an automobile engine. The cylinder has a volume of 250. cm 3 and the pressure of the gaseous octane is 3.50 atm in the hot engine (250 C). What is the pressure in the cylinder if the temperature of the engine is lowered to 125 C and the volume of the cylinder is changed to 500. cm 3? 2. A helium balloon needs to displace at least 1.00 x L of air. You fill the balloon with helium to a volume of 1.00 x L on the ground where the pressure is 745 mm Hg and the temperature is 20.0 C. When the balloon ascends to a height of 2 miles where the pressure is only 600. mm Hg and the temperature is -33 C, does the balloon still displace the required 1.00 x L of air? 3. When J.L Piccard made a stratosphere flight in a balloon, the balloon seemed to be only half-filled as it left the ground near Detroit. The gas temperature was 27 C, the pressure was 700. mm Hg, and the volume of gas in the balloon was 8.00 x 10 4 cubic feet. What was the volume at high altitude where the temperature was -3 C and the pressure was 400. mm Hg? 19

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21 MOLAR VOLUME OF A GAS Recall the trend in atomic size. Which is larger fluorine or chlorine? Since the (circle one) fluorine / chlorine molecule is larger, you would expect it to occupy a greater volume. Scientists in the 18th century believed this to be true and therefore disputed Avogadro when he made the following hypothesis in 1811: At constant temperature and pressure, equal volumes of gases contain equal number of molecules and therefore, equal number of moles. It was as if he were saying that two rooms of the same size (equal volumes) could be filled with the same number of objects no matter whether the objects were apples or marbles. But, since this theory is applied to gases, and not solids or liquids, it makes sense. From your everyday observation of gases and from the Kinetic Molecular Theory (KMT), we can conclude that the distance between gas molecules is very small/large compared to the size of the individual molecules. Therefore, doubling or tripling the size of a molecule still leaves a very large distance between them. The KMT states that these molecules are always/never moving in a set/random pattern and exert/do not exert attractive forces on each other. So on average, there is always going to be small/large expanses of space between the molecules. Avogadro demonstrated his theory through experiments showing that If the temperature is 0 C and the pressure is 1 atm standard temperature and pressure (STP) then 1 mole of ANY gas will occupy a volume of 22.4 L This is called the molar volume of a gas and can be used as a conversion factor. 1 mole gas at STP = 22.4L 1. Before release into the atmosphere, a weather balloon has a volume of L. How many moles of helium will occupy the balloon at ground level where the conditions are at STP? a. How many grams of helium are contained in the weather balloon? 2. A chemical reaction produces 98.0 ml of sulfur dioxide gas at STP. What is the mass, in grams, of the gas? 21

22 3. A BIC lighter contains 5.00 g of liquefied butane (C 4 H 10 ). What size balloon, in liters, at STP can be inflated with the butane lighter? a. If the lighter is at 20 C and 740. mm Hg, what is the volume, in liters, of the balloon? (Combined Gas Law) 4. Maggie Nesium blows up the lab as she decomposes 3.20 g of potassium chlorate. a. Write the balanced equation for the decomposition of potassium chlorate. b. Calculate how many grams of oxygen gas are produced. (Remember stoichiometry?) c. Calculate the volume, in liters, this gas will occupy at STP. (Molar volume) d. Calculate the volume, in liters, of this gas at 25 C and atm. (Combined Gas Law) 5. When making water, how many liters of oxygen are needed to completely react with 4 liters of hydrogen gas at STP? The reaction takes place at 20 C and 0.80 atm. 6. Using the Periodic Table, it is simple to find the mass of one mole of gas. Combining that with the fact that one mole of gas has a volume of 22.4 L at STP, the density of any gas at STP can be calculated. a. Determine the density of hydrogen gas at STP. 7. Determine the density of air at STP. Assume air to be 80% nitrogen and 20% oxygen. 22

23 IDEAL GAS EQUATION The four variables, P, V, T, and n, can be combined into a single equation, called the ideal gas equation. R = ideal gas constant = P units = V units = n = T units = EXAMPLES: 1. What P is exerted by moles of hydrogen gas in a 4.08 liter container at 35 0 C? 2. A gas sample occupies 8.77 L at 3.97 atm. What is the temperature given that there are 1.45 moles of gas in the sample? 3. What volume does 3.46 moles of a gas occupy at 45 0 C and 7.22 atm? 4. How many grams of carbon dioxide gas are there in a 4.1 L container at 34 0 C and 1.04 atm? 23

24 The Ideal Gas Equation can be rearranged to solve for molar mass and density. 1. MOLAR MASS (MM): 2. DENSITY (D): Examples 1. What is the molar mass of a gas if grams of the gas occupies a volume of 125 ml at 20 0 C and atm? 2. Find the density of NH 3 at 63 0 C and 705 mm Hg. 3. The density of a gas is 2.0 g/l at 1.5 atm and 27 0 C. Find the molar mass. 24

25 THE IDEAL GAS EQUATION An ideal gas is one in which the molecules have no volume and exert no attraction for one another. No ideal gases exist, but under certain conditions of temperature and pressure, real gasses approach ideal behavior. The four variables used to describe the state of a gas can be combined into a single expression: PV = nrt where R is a proportionality constant known as the ideal gas constant. The value and units of R depend upon the units of P, V, n, and T. In this class, we will use the value R = L-atm mole-k Since most of the time, chemists measure in grams, not moles, of a gas, the equation can be modified to include grams. Since moles = grams/molar mass, we can substitute g/mm for n in the ideal gas equation: PV = grt MM Using this equation, chemists can calculate the molar mass and density of a gas. Answer the following questions. Show all work. 1. At what temperature will 5.00 g of chlorine gas exert a pressure of 900. torr at a volume of 750. ml? 2. How many moles of oxygen will occupy a volume of 2.5 L at 1.2 atm and 25 0 C? x 10 4 g of helium are placed in a balloon. What is the volume of the balloon at 1.20 atm and 22 0 C? 25

26 4. What pressure (in kpa) will be exerted by 25 g of carbon dioxide at a temperature of 25 0 C and a volume of 500. ml? 5. What is the density of ammonia at 800. mm Hg and 25 0 C? 6. To find the volume of a flask, evacuate it so that it contains no gas. Next, 4.4 g of CO 2 is placed in the flask. On warming to 27 0 C, the gas exerts a pressure of 730 mm Hg. What is the volume of the flask? 7. If the density of a gas is 1.2 g/l at 14.4 psi and 20 0 C, what is its molar mass? 26

27 1. A metal cylinder contains 1 mole of nitrogen gas at STP. What will happen to the pressure if another mole of gas is added to the cylinder? 2. What three conditions influence the pressure exerted by a gas in a closed container? 3. Aerosol cans explode if heated. Why? 4. Methane is 23 times more powerful than carbon dioxide in contributing to global warming. Elsie the Cow produces, on average, 200 liters of methane gas per day! If Elsie, whose body temperature is 39 C, produces this volume of gas by digestion, what will the volume of that gas be when the temperature is 25 C? 5. Freon-12 (the common name for the compound CCl 2 F 2 ) was widely used in refrigeration systems, but has been replaced by other compounds that do not lead to the breakdown of the protective ozone in the upper atmosphere. Consider a 1.5 L sample of gaseous CCl 2 F 2 at a pressure of 740 torr. If the pressure is reduced to only 12 mm Hg in the upper atmosphere, what will be the new volume of the gas? 6. A bicycle tire is inflated to a pressure of 55 psi at 15 C. If the tire is heated to 35 C while riding in North Park, what is the new pressure in the tire? 7. If you release a helium-filled balloon, it soars upward and eventually pops. Explain this behavior. 27

28 8. The pressure exerted by a gas depends on the number of particles, the volume, and A. atmospheric pressure. B. temperature. C. the type of gas involved. D. the compressibility of the gas. 9. Which of the following would NOT cause an increase in the pressure of a gas? A. The container is made larger. B. More of the same gas is added to the container. C. The temperature is increased. D. Another gas is added to the container. 10. What is the effect on the volume of one mole of gas when the pressure is reduced to 1/3 of its original pressure? A. V remains constant. B. V is reduced to 1/3 of its original size C V triples. D. V doesn t change. 11. If the temperature is 294 K, what is the Celsius temperature? A. 273 o C B. 0 o C C. 567 o C D. 21 o C 12. Which of the following is NOT equivalent to standard pressure? A psi B. 1 kpa C. 760 torr. D. All of the above are equivalent to standard pressure. 13. Of the following relationships, which is inverse? A. P and T B. P and V C. P and n D. none of the above. 14. The definition of pressure is force/area. How can one increase the pressure exerted by an object? A. increase the area B. decrease the area C. decrease the force D. spread out the force 15. Calculate the density of neon at STP. 16. Radon, a radioactive gas formed naturally in the soil, can cause lung cancer. It can pose a hazard to humans by seeping into houses, and there is concern about this problem in many areas. A 1.5 mol sample of radon gas has a volume of 21.0 L at 33 C. What is the pressure of the gas? 28

29 17. One of the many problems encountered on the Apollo 13 space mission was a build-up of carbon dioxide gas. If the three astronauts on board exhaled a total of 18 kg of carbon dioxide at STP, what is the volume of gas that the astronauts produced in the space shuttle? 18. On a cold winter day, a person takes a breath of 0.50 L of air at 745 mm Hg and -6 C. What is the volume of this air in the lungs at 37 C and 1.00 atm pressure? 19. What will be the pressure (in kpa) exerted by the moles of carbon dioxide located at the top of a bottle of soda while the bottle is in the refrigerator at 5 o C? The volume of the gas is 50. ml. A. 2.0 kpa B. 3.6 x 10-5 kpa C. 2.0 x 10 2 kpa D kpa 20. The planet Krypton, the home of Superman, has an atmosphere different from earth. If the density of the atmosphere is 2.41 g/l at 10.6 psi and 33 o C, what is its molar mass? Name the elemental gas. A. 84 g/mole B g/mole C g/mole D. none of the above 29

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31 Man Your Stations! Now that you have learned 4 gas laws, you are going to apply your new knowledge to explain some common everyday occurrences. First, let s review. 1. What 4 factors are needed to define the state of a gas? 2. State Boyle s Law: 3. State Charles Law: 4. State Gay-Lussac s Law: 5. State Avogadro s Law: Station 1: TURNING A BALLOON INSIDE OUT Place a teaspoon of water in the flask and boil it on a hotplate until it is all almost gone. Using tongs remove the flask from the hotplate and immediately place a balloon over the top of the flask. Observe what happens; be patient! 1. What factors are set constant when the balloon, which acts as a lid, is placed on top of the flask? 2. Once the balloon is inside the flask, put the flask back on the hot plate and observe what happens (please remove the flask again before the balloon pops). 3. Explain what happened to the balloon in terms of the factors that affect gases. 4. What gas law could be used to explain this occurrence? 5. There are several methods that can be used to either put the balloon into the flask or remove it. However, all of them require a change in what variable? 31

32 Station 2: POPCORN Take a handful of popcorn from the bowl and enjoy as you work. Please be sure that no popcorn ends up on the floor! The reason why popcorn pops can easily be explained using the gas laws. Your mission is to figure out what is going on. I ll give you a hint inside each popcorn kernel, at the very core, is water. 1. Which factor changes as soon as you place the popcorn kernel in the microwave and turn it on? 2. How will this affect the water inside the kernel? 3. Using the factors that affect gases, explain why the popcorn pops. Be sure to state which factors remain constant until the very end. 4. What gas law is being addressed throughout most of the popping process? 5. What gas law takes over as the popcorn goes pop? Station 3: DEAD FISH Read the index card, and then answer the following questions. 1. Two of the four factors change as you come up from the very bottom of the ocean. Which ones are they and how do they change? 2. Explain why the fish died in terms of the factors that affect gases. 3. What 2 gas laws are illustrated in this problem? 32

33 Station 4: CARTESIAN DIVER Squeeze the sides of the bottle and observe what happens to the dropper inside the bottle. Release the pressure on the bottle, and again observe what happens to the diver. 1. What causes the diver to sink in the bottle? 2. Does the amount of air in the dropper change? Explain your answer. 3. What gas law does the Cartesian Diver illustrate? Station 5: HOW BUNNIES GROW INTO RABBITS Place a marshmallow bunny in the bell jar. Make sure the valve on the bottom of the pump is tight, then plug in the pump. Observe what happens to the bunny. To remove the bunny from the bell jar, loosen the screw to allow air to enter the jar and watch the bunny dance. 1. Which factor changes when the air is removed from the bell jar? 2. Why does the bunny experience instant growth? 3. What gas law is being illustrated here? 4. We just examined a relationship between pressure and volume, but other methods can be used to change the volume of a gas. What happens when a marshmallow is put into a cup of hot cocoa? Why? 5. What law could be used to explain this occurrence? 33

34 Station 6: CANDLE MAGIC This station is not a particular gas law but looks at the relationship between two of the four factors. Light the candle in the bowl. Carefully place the jar over the candle and observe. 1. What happened to the water when the candle went out? 2. For a hint as to what is occurring, balance the complete combustion reaction that occurs when a candle, C 30 H 62, burns. 3. By examining your equation, do the moles of gas increase or decrease after the reaction? (hint: compare the moles of O 2 to the moles of CO 2 ) 4. How does this affect the pressure of the gas inside the container? 5. Write this relationship as a law, just like you did for Boyle s and Charles. (Ignore the minimal temperature change.) 6. We still have not determined why the water rose in the jar. By comparing this demonstration to a barometer or open-end manometer, explain why. 34

35 What is the volume of one mole of gas at STP? It is easy to determine the mass of solids and liquids in the laboratory. One can simply use a balance to weigh the material. It is difficult, however, to find the mass of a gas. Chemists usually measure the volume of gas and then calculate its mass. In order to do this, the scientist must know the relationship between the molar mass and the molar volume of a gas. Avogadro s Hypothesis states that under the same conditions of temperature and pressure, equal volumes of gases contain equal number of molecules. Scientists then calculated the volume of one mole of gas at STP (273 K and 1 atm) and found it to be 22.4 liters. This is called the molar volume. The molar volume of a gas at STP is a convenient constant that we can use in stoichiometric calculations to determine the volume of gas that should be produced in a given reaction at a given temperature and pressure. It can also be used to determine the mass of a gas. Measure the volume of a gas and one can then calculate how many moles are present, and finally, the mass of that volume of gas. In this experiment, you will determine the value of this constant by using the data collected from a reaction between magnesium and HCl. You allow a known mass of magnesium to react with excess HCl and measure the volume of hydrogen gas evolved. Since molar volume is measured at STP, the volume of hydrogen gas collected must be corrected to standard conditions. Knowing the number of moles of hydrogen that should be produced and the volume of hydrogen corrected to STP, you can calculate the number of liters of gas per mole, i.e., the molar volume. OBJECTIVES: to determine an experimental molar volume for H 2. to study mass-volume relationships in chemical reactions. to measure the pressure of a dry gas, one that is collected over water. MATERIALS: Magnesium ribbon 800-mL beaker 6M HCl thermometer rubber stopper metric ruler 100-mL graduated cylinder thread barometer Got Gas? PROCEDURE: 1. Upon entering the lab, fill an 800-mL beaker almost completely full of tap water and allow it to stand so that the water will come to room temperature. 2. Record the temperature and pressure of the room. 3. Cut a piece of magnesium ribbon the length designated by your teacher; sand the ribbon. 4. Weigh the piece of magnesium to.001 g. It must weigh between grams. Record the mass in your data table. 5. Fold the magnesium ribbon so that it will fit into a 100-mL graduated cylinder. Tie a 15-cm piece of thread to it. 6. Slowly pour 25 ml of HCl into your 100-mL graduated cylinder. CAUTION: Dilute HCl will stain, cause burns, and irritate the lungs and eyes. Avoid contact and do not breathe the fumes. Rinse spills with plenty of water. 7. Tilt the graduated cylinder and fill it very slowly with the tap water from your beaker. Pour the water down the side of the tube so that the water and HCl do not mix. Fill the tube completely to the top so that no air remains inside. (Since hydrochloric acid has a density greater than water, the acid will remain at the bottom of the cylinder.) 8. With the tube completely filled with water, insert the magnesium ribbon about 3 to 4 cm into the cylinder. With the thread sticking out of the tube, quickly place the rubber stopper on top of the cylinder. (If you slant the stopper slightly, it will help to block the air hole that exists because of the spout on the tube.) 9. Using your fingers to cover up the stopper hole and spout hole, invert the cylinder into the beaker of water that has come to room temperature. Pour out some of the water in the beaker until it is only half full. 10. After the reaction is completed (all the magnesium will be gone), hold the graduated cylinder so that the water inside the cylinder is at the same level as the water in the beaker. 11. Read the volume of gas in the cylinder. Record in your data table. (Remember, you are reading an inverted cylinder!) 12. Empty the graduated cylinder and repeat the procedure steps 1 to 11 one more time. 35

36 Laboratory Assignment Use the following information to answer the questions. Show work, include units, and put your answers in the blanks. Molly B. Denim collects oxygen gas over water by decomposing 5.45 g potassium chlorate. The experiment took place in a lab with a temperature of 22 C and a pressure of 745 mm Hg. 1. Calculate the moles of potassium chlorate that were consumed in the reaction. 2. Calculate the number of moles of oxygen that should be produced in the reaction. 3. Using Dalton s Law of Partial Pressures, calculate the pressure of the dry oxygen gas. P Total = p 1 + p 2 + p Use the combined gas law to calculate the volume of O 2 that should have been collected at 0 C and 1 atm. 5. Determine the experimental molar volume. 6. In her second trial, Molly calculated a molar volume of 21.9 L/mol. Calculate the average molar volume. 7. Calculate the percent error of Molly s average molar volume. 36

37 Name Period Data Table Mass of Mg ribbon Trial 1 Trial 2 Room Temperature Room Pressure Volume of gas collected Water Vapor Pressure Calculations Table 1. Moles of Mg Trial 1 Trial 2 2. Moles of hydrogen gas 3. Pressure of dry H 2 gas 4. Volume of H 2 at STP 5. Experimental Molar Volume 6. Average Molar Volume 7. Percent Error 37

38 Laboratory Conclusion Perform the following calculations and write the answers in the spaces provided and in the Calculations Table. 1. Calculate the moles of potassium chlorate that were consumed in the reaction. 2. Calculate the number of moles of oxygen that should be produced in the reaction. 3. Using Dalton s Law of Partial Pressures, calculate the pressure of the dry oxygen gas. P Total = p 1 + p 2 + p Use the combined gas law to calculate the volume of O 2 that should have been collected at 0 C and 1 atm. 5. Determine the experimental molar volume. 6. In her second trial, Molly calculated a molar volume of 21.9 L/mol. Calculate the average molar volume. 7. Calculate the percent error of Molly s average molar volume. 38

39 WHAT do I need to know?? Unit 12: The Gas Laws PART I: multiple choice o Kinetic Molecular Theory gases o Definition of pressure how atmospheric pressure changes with elevation how atmospheric pressure effects melting and boiling points o Variables that affect gases and their relationship to each other and to the behavior of the gas P, V, T, n inverse or direct relationships o The Gas Laws Boyle s Charles Gay-Lussac s Combined Ideal Be able to solve problems there are several imbedded in the multiple choice section o Temperature conversions 0 C K o F.A.R.T. lab o Man Your Stations lab 39

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