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UNIT 5 NOTES Objectives: Physical Characteristics of Gases Use the kinetic molecular theory (KMT) to explain how certain physical properties of ideal gases differ from real gases Describe the conditions under which real gases deviate from "ideal" behavior Explain the five postulates of the kinetic molecular theory. Define pressure and standard pressure in terms of force and explain how pressure is measured Use the gas laws to express simple mathematical relationships between the pressure, temperature, volume and quantity of gases Relate KMT to the individual gas laws Molecular Composition of Gases Develop a relationship between the volume, mass, and number of particles of a gas Perform calculations using the ideal gas law, specifically being able to solve for pressure, volume, temperature, or number of moles Perform stoichiometric calculations using the relationship between volume and moles both at STP and non STP conditions Demonstrate the relationship between the mass of gas particles and their rate of effusion Vocabulary: Kinetic molecular theory Gas pressure Gas temperature Ideal gas Ideal gas constant (R) Standard temperature and pressure (STP) Standard atmosphere (atm) Kelvin (K) Pascal (Pa) Dalton’s law of partial pressures Effusion Diffusion Graham’s law Millimeter of mercury/Torr (mmHg) Kinetic energy Avogadro’s Law Boyle’s Law Charles’ Law Gay-Lussac’s Law rate Properties of Gases To understand why a gas behaves as it does, we need to be able to picture what happens to gas particles as conditions (like temperature and pressure) change. The Kinetic Molecular Theory (the theory of moving molecules) provides this picture. It is summarized below. 1. Gases consist of large numbers of molecules that are in constant, random motion. 2. The individual molecules of a gas essentially have zero volume compared to the total volume of the gas. 3. Attractive and repulsive forces between gases are negligible. 4. The particles move in straight lines. 5. The average energy of the particles is directly proportional to absolute temperature (Kelvin). At any given temperature all gases have the same kinetic energy. If a gas behave in accordance with the above 5 postulates it is referred to as ideal. Real gases deviate from which of the postulates? Scientists use the “Gas Laws” to describe what will happen to gases under certain circumstances. There are several gas laws that we will discuss, but they are all derived from the same basic gas equation: PV = nRT As you can see, this equation has many variables. Each one will have a different affect on how a gas behaves. The Variables: 1. Temperature (T) a. Defined as the average kinetic energy that the particles of the gas possess. b. The number used to represent temperature depends on the scale being used. There are 3 scales: i. Fahrenheit, Celsius, and Kelvin. 1. Kelvin is considered the absolute temperature scale because 0 K is called absolute 0 and is theoretically the lowest possible temperature and at this temperature the kinetic energy is zero. c. Most thermometers we use record Celsius and/or Fahrenheit. Very few thermometers record Kelvin. To determine Kelvin we must do a conversion. i. To convert temperature into Kelvin, you must start with a temperature in Celsius! Use: TK = TC + 273 Example: 100oC = _________K ii. All temperature in gas law problems must be in Kelvin! 2. Pressure (P) a. Defined as the force exerted on an object divided by the area over which the force is exerted. i. Exerting more force increases the pressure ii. Exerting a force over a smaller area increases the pressure. iii. We use air under pressure (compressed air) for many everyday uses like filling our car tires with air and to fill SCUBA diving tanks. iv. Pressure is the result of billions of gas particles moving and colliding against the container in which the gas is found. b. Units of Pressure – There are many units used to measure pressure. The SI unit for pressure is atm mm Hg torr Atmospheres 1 760 c. d. 3. 760 psi pounds per square inch 14.7 kPa kilopascals Pa Pascals 101.325 1.01325x105 Standard Temperature and Pressure – STP i. This is the temperature and pressure which is universally known all around the world to be 273K and 101.325 kPa ii. Sometimes STP uses the pressure unit “atmosphere” instead, in which case STP is 273K and 1 atm. List three ways you can increase the pressure i. _______________________________________ ii. _______________________________________ iii. _______________________________________ Volume (V) a. Defined as the amount of space the gas takes up. b. Units of Volume – There are many units used to measure volume. The SI unit for volume 1 dm3 = 1 liter as well as: 3 cm quarts Liter 1 1000 1.057 c. 4. Remember that gas is the one phase of matter that does not have a definite volume. The gas laws we will be discussing will describe how/when the volume of a gas changes. Moles (n) a. The term mole is used to describe a certain number of gas particles. b. Remember that 1 mole = 6.02 x 1023 particles c. Avogadro discovered that 1 mole of any gas (provided it is at 1 atm of pressure and 273K) has exactly 6.02 x 1023 particles and 22.4 L = 1 mol. d. In chemistry, this is useful because you can relate the number of particles in 1 mole of any gas to its atomic mass. For example: 1 mole oxygen = 6.02 x 1023 particles = 32 grams= 22.4 Liters at STP The Volume-Amount of Gas Relationship: Avogadro’s Law As the amount of gas increases in an elastic container, the volume will increase (pressure and temperature held constant. As the amount of gas decreases in an elastic container, the volume will decrease (pressure and temperature held constant. Why does this happen? Avogadro’s Law V 1 = V2 n 1 n2 or V1n2 = V2n1 Sample Problem: 1. 5.00 L of a gas is known to contain 0.965 mol. If the amount of gas is increased to 1.80 mol what new volume will result (at an unchanged temperature and pressure)? V1n2 = V2n1 5.00 L(1.80 mol) = V2(0.965 mol) V2 = 9.33 L 2. 13.1 mol of a gas is in a 4.37 L container at constant pressure. The volume of the container decreases to 2.72 L how many moles of The Pressure – Volume Relationship: Boyle’s Law The first person to investigate the relationship between the pressure of a gas and its volume was the British chemist Robert Boyle (1627 – 1691). As the pressure of a gas increases, its volume decreases and as the pressure of a gas decreases, the volume increases (temperature held constant). As the volume of a gas increases, its pressure decreases and as the volume of a gas decreases, its pressure increases (temperature held constant). Why does this happen? Boyle’s Law P1V1 = P2V2 or P1 = P2 V2 V1 Sample Problem: 1. The gas in a balloon has a volume of 4L when it is at 100kPa of pressure. The balloon is released into the atmosphere, where it expands to a volume of 8L. What is the new pressure on the balloon? P1 V1 = P2 V2 (100 kPa)(4 L) = (P2) (8 L) 100kPa (4L) = P2 8L P2 = 50 kPa 2. Consider a 1.53 L sample of a gaseous SO2 at a pressure of 5.6 x 103 Pa. If the pressure is changed to 1.5 x 104 Pa at constant temperature, what will be the new volume of gas? The Temperature – Volume Relationship : Charles’ Law Jacques Charles (1746-1823) was the scientist who developed the scientific law that relates temperature of a gas to its volume. As the temperature of a gas increases, its volume increases and as the temperature of a gas decreases, its volume decreases (pressure is held constant). As the volume of a gas increases, its temperature increases and as the volume of a gas decreases, its temperature decreases (pressure is held constant). Why does this happen? Charles’ Law V1 = V2 T1 T2 or V1T2 = V2T1 Sample Problem: 1. Gas in a balloon occupies 2.5 L at 300K. At what temperature will the balloon expand to 7.5 L? V1 V2 T1 T2 T2 T2 V2T1 V1 7.5 300 T2 900 K 2.5 2. A gas with a volume of 600. mL has a temperature of 30°C. At constant pressure the gas is heated until the gas expands to 1,200 mL. What is the new temperature of the gas if the pressure remains constant? The Temp. – Pressure Relationship : Gay-Lussac’s Law Joseph Gay-Lussac explored the relationship between the temperature of a gas and its pressure. As the pressure of a gas increases, its temperature increases and as the pressure of a gas decreases, its temperature decreases (volume held constant). As the temperature of a gas increases, its pressure increases and as the temperature of a gas decreases, its pressure decreases (volume held constant). Why does this happen? Gay-Lussac’s Law P1 = P2 T1 T2 or P1T2 = P2T1 Sample Problem: 1. The pressure of a gas in a tank is 4.20 atm at 44oC. If the temperature rises to 67oC, what will be the gas pressure in the tank? P1 P2 T1 T2 P2 P1T2 T1 4.20 atm (340 K ) 317 P2 4.50 atm 2. Ten liters of a gas is found to exert 97.0 kPa at 25.0°C. What would be the required temperature (in Celsius) to change the pressure to standard pressure? Combined Gas Law There are times when temperature, volume, and pressure are all affected when the conditions of a gas change. In these instances, we must combine Boyle’s Law, Charles’ Law, and Gay-Lussac’s Law into what is known as the Combined Gas Law. Knowing the combined gas laws gives the individual gas laws by holding one of the variables constant. Combined Gas Law P1V1 = P2V2 T1 T2 Sample Problem: 1. A helium balloon with a volume of 410 mL is cooled from 27oC to –27oC. The pressure on the gas is reduced from 100kPa to 25 kPa. What is the volume of the gas at the lower temperature and pressure? 2. A gas has a volume of 800.0 mL at minus 23.00 °C and 300.0 torr. What would the volume of the gas be at 227.0 °C and 600.0 torr of pressure? 3. 500.0 liters of a gas are prepared at 700.0 mm Hg and 200.0 °C. The gas is placed into a tank under high pressure. When the tank cools to 20.0 °C, the pressure of the gas is 30.0 atm. What is the volume of the gas? (Don’t forget to convert mm Hg to atm). 4. What is the final volume of a 400.0 mL gas sample that is subjected to a temperature change from 22.0 °C to 30.0 °C and a pressure change from 760.0 mm Hg to 360.0 mm Hg? 5. What is the volume of gas at 2.00 atm and 200.0 K if its original volume was 300.0 L at 0.250 atm and 400.0 K? 6. At conditions of 785.0 torr of pressure and 15.0 °C temperature, a gas occupies a volume of 45.5 mL. What will be the volume of the same gas at 745.0 torr and 30.0 °C? 7. A gas occupies a volume of 34.2 mL at a temperature of 15.0 °C and a pressure of 800.0 torr. What will be the volume of this gas at standard conditions? Molar Volume and Reactions of Gases As you solve the following problems keep in mind Avogadro’s law, which states that equal volumes of all gases at the same temperature and pressure contain the same number of molecules. From Avogadro’s law, it follows that all gases have equal molar volumes if they are measured at the same temperature and pressure. The molar volume is the volume occupied by one mole of a substance. At standard temperature and pressure (STP―0°C and 1 atm), the molar volume of any gas sample is 22.4 L. 1. What volume would be occupied by 2.0 mol nitrogen, N2, gas at 0°C and 1 atm? Volume = ___________________ 2. What volume would be occupied by 88.0 g of gaseous carbon monoxide, CO, at STP? Volume = ___________________ 3. Tanks of gaseous propane are used for cooking and heating. When propane (C3H8) burns (using oxygen from the air), the products of the reaction are carbon dioxide and water vapor. a. Write a balanced equation for this reaction. ___________________________________________________________________________ b. At constant temperature and pressure, how many liters of oxygen would be needed to completely combust 0.350 L of propane? Volume = ___________________ c. Continuing from 3b, how many liters of water vapor would be produced by the reaction of 0.350 L of propane? Volume = ___________________ 4. Hydrogen chloride gas can be produced by a reaction between hydrogen gas and chlorine gas. a. Write a balanced equation for this reaction. ________________________________________________________________________ b. At constant temperature and pressure, how many liters of hydrogen are needed to produce 1.75 L of hydrogen chloride? Volume = ___________________ c. Continuing from 4b, how many moles of chlorine would be needed to react with 8.65 mol of hydrogen? Volume = ___________________ Ideal Gas Law PV = nRT Volume, pressure, temperature, and amount of gas can be compared by using a conversion factor known as the ideal gas constant (R). While the term constant implies not changing, the value of the gas constant is dependent on the unit used to measure the pressure. Gas Constant Values If the Units of pressure is: The Numerical Value of R Is (with units): atm kPa mmHg (Torr) PV = nRT R = PV nT 0.0821 L-atm/mole-K 8.314 L-kPa/mole-K 62.4 L-Torr/mole-K To solve in units of pressure: atm 1atm(22.4L) 1 mol(273K) *THE CONSTANT (R) THAT IS USED IS BASED ON THE UNITS OF PRESSURE = 0.0821 atm.L mol.K 1. If I have 4 moles of a gas at a pressure of 5.6 atm and a volume of 12 liters, what is the temperature? 2. If I have an unknown quantity of gas at a pressure of 120 kPa, a volume of 31 liters, and a temperature of 87 0C, how many moles of gas do I have? 3. If I contain 3 moles of gas in a container with a volume of 60 liters and at a temperature of 400 K, what is the pressure in mmHg inside the container? 4. If I have 7.7 moles of gas at a pressure of 0.09 atm and at a temperature of 56 0C, what is the volume of the container that the gas is in? 5. If I have 17 moles of gas at a temperature of 67 0C, and a volume of 88.89 liters, what is the pressure in kPa of the gas? 6. If I have an unknown quantity of gas at a pressure of 380 mmHg, a volume of 25 liters, and a temperature of 300 K, how many moles of gas do I have? The Ideal and Combined Gas Laws PV = nRT or P1V1 = P2V2 T1 T2 Use your knowledge of the ideal and combined gas laws to solve the following problems. If it involves moles or grams, it must be PV = nRT 1. If four moles of a gas at a pressure of 5.4 atmospheres have a volume of 120 liters, what is the temperature? 2. If I initially have a gas with a pressure of 84 kPa and a temperature of 35°C and I heat it an additional 230 degrees, what will the new pressure be? Assume the volume of the container is constant. 3. My car has an internal volume of 2600 liters. If the sun heats my car from a temperature of 20°C to a temperature of 55°C, what will the pressure inside my car be? Assume the pressure was initially 760 mm Hg. 4. How many moles of gas are in my car in problem #3? 5. A toy balloon filled with air has an internal pressure of 1.25 atm and a volume of 2.50 L. If I take the balloon to the bottom of the ocean where the pressure is 95 atmospheres, what will the new volume of the balloon be? How many moles of gas does the balloon hold? (Assume T = 285 K) Gas Stoichiometry Practice The amount of gas can be determined from the amount of another substance using stoichiometery. If the reaction takes place at STP, the molar volume relationship of 1 mol occupies 22.4 L can be used, however, if the reaction takes place at an conditions other than STP, the ideal gas equation must be used to solve for the mols of gas. 1. The industrial production of ammonia proceeds according to the following Equation. N2(g) +H2(g) NH3(g) a. What volume of nitrogen at STP is needed to react with 57.0 mL of hydrogen measured at STP? b. What volume ofNH3 at STP can be produced from the complete reaction of 6.39 x 104 L of hydrogen? c. If 20.0 mol of nitrogen is available, what volume ofNH3 at STP can be produced? d. What volume of H2 at STP will be needed to produce 800. L of ammonia, measured at 55°C and 0.900 atm? 2. Propane burns according to the following equation. C3H8(g) + O2(g) CO2(g) + H2O(g) a. What volume of water vapor measured at 250°C and 1.00 atm is produced when 3.0 L of propane at STP is burned? b. What volume of oxygen at 20.°C and 102.6 kPa is used if 640. L of CO2 is produced? The CO2 is also measured at 20.°C and 102.6 kPa. c. If 465 ml. of oxygen at SIP is used in the reaction, what volume of CO2, measured at 37°C and 0.973 atm, is produced? d. When 2.50 L of C3H8 at SIP burns, what total volume of gaseous products is formed? The volume of the products is measured at 175°C and 1.14 atm. 3. Silicon tetrafluoride gas can be produced by the action of HF on silica according to the following equation. SiO2(s) + HF(g) SiF4(g) + H2O(l) 1.00 L ofHF gas under pressure at 3.48 atm and a temperature of 25°C reacts completely with SiO2 to form SiF4. What volume of SiF4 , measured at 15°C and 0.940 atm, is produced by this reaction? 4. One method used in the eighteenth century to generate hydrogen was to pass steam through red-hot steel tubes. The following reaction takes place. Fe(s) + H2O(g) Fe3O4(s) + H2(g) a. What volume of hydrogen at SIP can be produced by the reaction of 6.28 g of iron? b. What mass of iron will react with 500. L of steam at 250.°C and 1.00 atm pressure? c. If 285 g ofFe3O4 are formed, what volume of hydrogen, measured at 20.°C and 1.06 atm, is produced? 5. Sodium reacts vigorously with water to produce hydrogen and sodium hydroxide according to the following equation. Na(s) + H2O(l) NaOH(aq) + H2(g) If 0.027 g of sodium reacts with excess water, what volume of hydrogen at STP is formed? 6. In one method of producing aluminum chloride, HCl gas is passed over aluminum and the following reaction takes place. Al(s) + HCl(g) AlCl3(g) + H2(g) a. What mass of Al should be on hand in order to produce 6.0 x 103 kg of AlCl3? b. What volume of compressed HCI at 4.71 atm and a temperature of 43°C should be on hand at the same time? 7. It is possible to generate chlorine gas by dripping concentrated HCl solution onto solid potassium permanganate according to the following equation. KMnO4(aq) + HCl (aq) KCl(aq) + MnCl2(aq) + H2O(l) + Cl2(g) If excess HC1 is dripped onto 15.0 g of KMnO4 ,what volume of C12 will be produced? The Cl2 is measured at 15°C and 0.959 atm. 8. One of the reactions in the Solvay process is used to make sodium hydrogen carbonate. It occurs when carbon dioxide and ammonia are passed through concentrated salt brine. The following equation represents the reaction. NaCl(aq) + H2O(l) + CO2(g) + NH3(g) NaHCO3(s) + NH4Cl(aq) a. What volume ofNH3 at 25°C and 1.00 atm will be required if38 L of CO2, measured under the same conditions, react to form NaHCO3? b. What mass of NaHCO3 can be formed when the gases in (a) react with NaCl? c. If this reaction forms 46.0 kg of NaHCO3, what volume of NH3, measured at STP, reacted? Dalton’s Law of Partial Pressure: Dalton stated that the sum of the pressures of individual gases in the same container is equal to the pressure of the container. Ptot=P1+P2+P3… Effusion, Diffusion, and Graham’s Law Gases are a fluid substance in constant random motion. It is possible to measure the velocity of a gas based on its kinetic energy. Before beginning let’s review a few concepts. What exactly is temperature a measurement of?________________________________ Why is it important to include the word "average" in your answer? _________________ What two factors does an object's kinetic energy depend on? _________ and ________ What specifically is the equation for kinetic energy? _____________________________ Which would increase the kinetic energy of an object more: doubling the object's mass or doubling the objects velocity? _________ Explain: ____________________________ Define diffusion___________________________________________________________ Define effusion___________________________________________________________ State Graham's Law as an equation for two gases (A and B) at the same temp: Based on Graham’s Law the velocity of a gas varies inversely with its mass in kilograms. Consider two gases, He and O2, at the same temperature... Which particles would have greater average kinetic energy? ____________ Which particles are heavier? _____________ Which particles would have greater velocity? _____________ Which gas would diffuse across the room faster? ____________ Two gas samples, one H2 and one CO2, are such that their particles have the same velocity... Which gas molecules have the greater average kinetic energy? __________________________ Which gas is at the higher temperature? _______ Explain: ____________________________ For the following questions, use the Graham's Law equation. Show all work. At a certain temperature, O2 molecules move with an average velocity of 345 mph. At that same temperature, what would be the average velocity of a) He atoms? b) CO2 molecules? Ans: a) ________ b) ________ At a certain temperature, CH4 molecules move with an average velocity of 187 m/sec. At that same temp, gas X particles have an average velocity of 141 m/sec. a) Is gas X heavier or lighter than CH4? b) What is the molecular weight of gas X? c) What is a possible identity of gas X? Ans: a) ________ b) ________ c) ________ Molar Mass of a Gas In this lab activity you will collect a sample of CO2 gas over water from the unfinished/unbalanced reactions listed below: Reaction 1: CaCO3 (s) + Reaction 2: H2CO3 (aq) H2O (l) We will use antacid tablets (Alka Seltzer®) to generate the gas. Notice that carbonic acid formed in the first reaction decomposes in the second reaction to generate the CO2. From your lab data you will calculate the molar mass of carbon dioxide gas. Procedure Materials: Pneumatic trough Gas collecting bottle Glass plate Glass tubing 250 ml Flask Stopper Rubber tubing 100 ml Graduated cylinder Antacid tablet Balance Part 1 1. Fill your pneumatic trough with tap water, making sure the water level rises above the tray. 2. Fill your gas collecting bottle to the rim with tap water. Pour the water into a 100ml graduated cylinder. To get the total volume of the bottle, you may need to repeat this several times, emptying the cylinder between readings. Be sure to keep track of the total volume of water that was in the bottle. Record your volume in the data table below. 3. Refill the bottle with tap water so that it overflows a bit. Cover the bottle with the glass plate, making sure no air (little air) has been trapped; water may squirt out the sides as you do so. Holding onto the glass cover, invert the bottle and submerge it in the trough. Once submerged, remove the glass and slide the bottle onto the tray. You may need to lift the bottle a bit to do so, but BE CAREFUL NOT TO PULL THE BOTTLE COMPLETELY OUT OF THE WATER (TROUGH) OR YOU WILL HAVE TO RESTART. Position the bottle so that its neck is over the hole in your tray. 4. Make sure your glass bend tightly fits into the rubber tubing so no gas escapes. Insert the glass tubing halfway into the bottle through the hole in the tray. 5. Make sure the other end of the rubber tubing fits tightly onto the arm of the flask. Part 2 1. Add tap water to a 250 mL Erlenmeyer flask with arm until it is approximately half full. 2. Find the mass of the flask and tap water. Record your data in the table below. 3. Find the mass of ¼ of an antacid tablet. Record your data in the table below. Part 3 1. Add the antacid tablet to the flask and quickly cap the flask with the rubber stopper. 2. Once the flask stops fizzing, and the liquid level in the bottle stops moving, remove the glass bend from the bottle. BE CAREFUL NOT TO TIP IT OVER!! 3. Slide the bottle off of the tray and slowly raise or lower it so the liquid level inside matches the level in the trough. Be careful not to pull it out of the trough completely!! 4. Once the water levels match, cover the mouth of the bottle with the glass plate and remove the bottle from the trough. 5. Pour remaining water from the bottle into the graduated cylinder, and record your volume. 6. Find the mass of the flask and its remaining contents. Record the mass. 7. Record the room temperature and atmospheric pressure. Data Volume Data: Measurement Total volume of bottle (mL) Volume of remaining water (mL) Volume of gas (mL) Mass Data Measurement Mass of Flask + Water (g) Mass of antacid tablet (g) Total mass of reactants (g) Mass of flask + contents remaining (g) Mass of gas (g) Pressure Data Atmospheric Pressure (mmHg) Partial Pressure of H2O (mmHg) Partial Pressure of CO2 (mmHg) Trial 1 Trial 2 Trial 1 Trial 2 Temperature = ___________ºC vvvvvvvvvvvv___________K Calculations 1. Calculate the molar mass of CO2 from your lab data in the space below. Show all work!!!! 2. What is your percent error? (Hint: determine the accepted value for molar mass.) 3. What is the density of CO2 under your laboratory conditions? UNIT 5 REVIEW 1. Gases are distinguished from other states of matter by which of the following? a. Expansion b. Compressibility c. Homogeneity d. Space between atoms/molecules e. Constant and random motion f. All of the above 2. Which statement is not part of the Kinetic Molecular Theory of Gases? a. Gas atoms/molecules travel in straight-line motion and obey Newton’s Laws b. Collision between gas atoms/molecules are perfectly elastic (no energy gained or lost) c. Gases are composed of atoms/molecules of very small volume d. There are no attractive or repulsive forces between gas atoms/molecules 3. Define the differences between an Ideal Gas and a Real Gas. Why does this distinction need to be made? 4. Gases behave most ideally when? 5. The average kinetic energy of the particles of a gas: a. Is not affected by the temperature of the gas b. Increases as the temperature of the gas increases c. Decreases as the temperature of the gas increases d. Is equal to the total thermal energy absorbed by the substance 6. Write the formula for kinetic energy below, and define the terms involved: 7. Which of the following gases has the highest average kinetic energy at the same temperature and pressure? a. Oxygen b. Helium c. Hydrogen d. Nitrogen e. Neon 8. Which of the following gases has the highest average velocity at the same temperature and pressure? a. Oxygen b. Helium c. Hydrogen d. Nitrogen e. Neon 9. Consider a 2.47 L sample of gaseous SO2 at a pressure of 4.21 kPa. If the pressure is changed to 19 kPa at a constant temperature, what will be the new volume of the gas (in L)? 10. 3.00 L of a gas is known to contain 0.840 mol of molecules. If the amount of gas is increased to 1.60 mol, what new volume will result (in L), assuming an unchanged temperature and pressure? 11. What volume (in L) will 1.50 mol of oxygen (O2) occupy at -15°C and 1.8 atm? 12. The following reaction occurs at STP: C2H4 (g) + 3O2 (g) 2CO2 (g) + 2H2O (g) How many liters of CO2 gas are produced when 305.2 grams of C2H4 are consumed? 13. A sample of gas at 12.0°C occupies 400 mL under a pressure of 820 torr. To decrease the volume of this gas to 300 mL and decrease its temperature to 8.00°C, what pressure (in atm) must be achieved? 14. 8.00 L of a gas is found to exert 3.00 kPa of pressure at 20.0°C. Assuming constant volume, what would be the required temperature (in Celsius) to change the pressure to standard pressure? 15. A balloon contains 0.1 moles of oxygen and 0.4 moles of nitrogen. If the balloon is at STP what is the partial pressure of the nitrogen? 16. A gas with a volume of 400 mL has a temperature of 20.0°C. The gas is heated at constant pressure, and it expands to a volume of 1000 mL. What is the temperature (in K) of the gas after being heated? 17. 4.00 L of a gas is collected at 25.0°C and 800.0 mm Hg. What is the volume (in L) of this gas at STP? 18. CO2 gas is contained in a 2.00 L flask at 4.00°C, 109 kPa. How many molecules of CO2 are present? 19. 3.00 moles of bromine gas (Br2) has a temperature of -20.0°C at 105 kPa. What is the density (in g/L) of this gas? 20. What is the density (in g/L) of ethane gas, C2H6, at 810. mm Hg and 14.0°C? 21. KClO3 decomposes by the following reaction: 2 KClO3 (s) 2KCl(s) + 3O2(g) The O2 produced was collected by the displacement of water at 22°C at a total pressure of 760 mmHg. The volume of gas collected was 1.20 liters, and the vapor pressure of water at 22°C is 21 mmHg. Calculate the partial pressure (in atm) of O2 in the gas collected: