Honors Chemistry - Unit 11

Transcription

1 Honors Chemistry - Unit 11 Chapters 10 & 11 Gases, Gas Laws, and Gas Stoichiometry Vocabulary Due: UT Quest(s): VOCABULARY: Quizzes: Test Date: Ideal gas standard atmospheric pressure standard temperature STP diffusion effusion molar volume CONSTANTS/FORMULAS: Boyle s law Charles law Gay-Lussac s law Combined gas law Dalton s law Graham s law Ideal gas law Molar volume Molecular Mass Determination (of a gas) Unit 11 Packet - Page 1 of 14 OBJECTIVES: Memorize the values for STP. Memorize and be able to apply the gas laws: Boyle s, Charles, Dalton s law of partial pressure, Combined gas law, Gay-Lussac s, and Graham s. Be able to use molar volume of a gas at STP in problems. Be able to calculate gas density at STP. Memorize and be able to apply the ideal gas law. Memorize the gas constant R =.0821 L-atm/mol-K. Be able to do problems involving gas stoichiometry (at STP and other conditions). STP: Standard temperature and Pressure; 0 0 C(273K) and 1 atm Pressure conversions (all of these values are equal to each other and any two can be set-up as a ratio to be used in a factor label problem!) 1 atm = 760 torr = 760 mm Hg = kpa = 1.01 x 10 5 Pa PRESSURE CONVERSION PROBLEMS (Show work on your own paper and attach to packet). 1) 560 torr = kpa 2) 72.0 kpa = mm of Hg 3) 435 mm of Hg = atm 4) 1100 torr = mm Hg 5) 1.6 atm = kpa 6) 48.6 kpa = atm 7) 0.8 atm = torr 8) 82 torr = kpa 9) 760 mm Hg = atm 10) 2.0 kpa = torr 11) 485 mm Hg = kpa Place these eight pressure measurements in order from largest amount of pressure to smallest amount of pressure: A. 1.1 atm B. 59 kpa C. 14 mm Hg D kpa E mm Hg F kpa G. 9.2 atm H. 436 torr

2 Unit 11 Packet - Page 2 of 14 Unit 11 BELLWORK SET Gases, Gas Laws, and Gas Stoichiometry Review: 1. Draw the dot diagram for Bismuth. 2. Draw the Lewis Structure for the following and identify the VSEPR Shape: A. Carbonate ion B. silicon disulfide Current Unit Material 3. A sample of diborane gas (B2H6) a substance that bursts into flame when exposed to air, has a pressure of 345 torr at a temperature of 15 0 C and a volume of 3.48 L. If conditions are changed so that the temperature is 36 0 C and the pressure is 268 torr, what will be the volume of the sample? 4. The density of a gas was measured at 1.30 atm and 47 0 C and found to be 1.95 g/l. Calculate the molar mass of the gas. 5. Mixtures of helium and oxygen are used in scuba diving tanks to help prevent the bend. For a particular dive, 46 L of O2 at 25 0 C and 1.0 atm and 12 L He at 25 0 C and 1.0 atm were pumped into a tank with a volume of 5.0 L. Calculate the partial pressure of each gas and the total pressure in the tank at 25 0 C. Hint: Label all of the variables and calculate each tank (oxygen and helium) as separate problems! 6. Calculate the volume of oxygen gas at 85 0 C and 789 torr, required for the complete combustion of 45 g of octane (C8H18).

3 Unit 11 - Gas Laws Notes Gases: 4 measurable quantities: Unit 11 Packet - Page 3 of 14 Volume 1 ml = 1 cm 3 Pressure Temperature # of moles Variables? STP? **GAS LAWS Must be Memorized!!**! Boyle s law - P1 V1 = P2 V2 (V varies inversely with P) Graph of P vs V? inversely? Example. If 1.0 L of a gas at 1.2 atm is allowed to expand to 5.0 L, what is the new pressure? You try: A sample of O2 gas has a volume of 150 ml when its pressure is 720 mm Hg, what will the volume be if the pressure is increased to 750 mm Hg? (Answer: 144 ml) Charles law: V1/T1 = V2/T2 (V varies directly with the Kelvin temperature) T must be in Kelvin Graph of V vs T? Directly? Why T in K? Example: A helium filled balloon has a volume of 2.75 L at C. The volume of the balloon decreases to 2.46 L after it was placed outside on a cold day. What is the outside temperature? You try: A sample of neon occupies a volume of 752 ml at 25 0 C. What volume will the gas occupy at C? (Answer: 815 ml)

4 Gay-Lussac s Law: P1/T1 = P2/T2 Like which other law? Unit 11 Packet - Page 4 of 14 (P varies directly with Kelvin Temperature) You Try: Before a trip from Raleigh to NY, the pressure in a tire is 1.8 atm at 20 0 C. At the end of the trip the pressure gauge reads 1444 mm Hg. What is the new temperature in the tire? (Answer: 309 K (or 36 0 C)) Combined Gas Law: P1V1 = P2V2 T1 T2 T must be in K!!! Example: If ml of a gas at 25 0 C and 1.3 atm is cooled to C and 780 mm Hg. What is the new volume? You try: a 700. ml gas sample at STP is compressed to a volume of 200. ml, and the temperature is increased to C. What is the new pressure of the gas in kpa? (Answer: 394 kpa)

5 Unit 11 Packet - Page 5 of 14 Dalton s law of Partial Pressures: pressure of each gas in a mixture is called the partial pressure of that gas. Total Pressure = sum of partial pressures PT = P1 + P2 + P3. etc. Special case: gases collected by water displacement see diagram on board! Use formula PT = Patm = Pg + PH2O Total atmosphere gas water Patm from lab or is given: PH2O = Table page 899 Example: Oxygen is collected by water displacement. The barometric pressure and temperature are 84.5 kpa and C. What is the partial pressure of O2? You try: A gas is collected over water and the atmospheric pressure is kpa at 50 C. What is the partial pressure of the gas? Graham s Law Review: Effusion? Diffusion? **Rate of effusion and diffusion varies inversely with the mass of the gas.** Rate of gas A = Square Root Mass of gas B Rate of gas B Mass of gas A Example: Compare the rates of effusion of Hydrogen and oxygen at STP. (Answer: 88.8 kpa) You try: Compare the rates of effusion of Fluorine and chlorine (Answer: Fluorine is 1.4 x faster than chlorine)

6 Ideal Gas Law Notes PV = nrt Must use the following units with the ideal gas law! P = atm V = L T = K n = moles Unit 11 Packet - Page 6 of 14 R = gas constant L-atm/mol- K (memorize) Example: What is the pressure exerted by a 12.0 g sample of Nitrogen gas (N2) in a 10.0 L container at 25 0 C? Practice Ideal Gas Law Worksheet: 1 4 (page 8 in packet) Gas Stoichiometry Molar Volume - 1 mol of any gas at STP has a volume of 22.4 L 1 mol = conversion factor (only to be used at STP) 22.4 L Example: What is the mass of 98.0 ml of sulfur dioxide at STP? (work together on board copy into your notes!) You try: What is the mass of 1.33 x 10 4 ml of O2 at STP? Gas Density: D =? At STP : D = Molar mass Molar volume g/mol (PT) 22.4 L/mol Ex: What is the density of CO2 at STP? You try: What is the molar mass of a gas with a density of 1.56 g/l?

7 Molecular Mass Determination from Ideal Gas Law: (see derivation on board) D = PM density = pressure (atm) x molar mass/ constant x temp (K) RT Examples & practice on ideal gas law worksheet! Stoichiometry of Gases Can do L-L conversions (just like mol-mol) with 2 gases and an equation Example: Given: C3H8 (g) + 5 O2 (g) à 3 CO2 (g) + 4 H2O (l) How many L of O2 are required to react with 0.35 L of propane? **You may have to use molar volume!** Unit 11 Packet - Page 7 of 14 **You may have to use the ideal gas law!** Example 2: Given CaCO3 (s) à CaO (s) + CO2 (g) How many grams of calcium carbonate must decompose to from 5.0 L of CO2 at STP? You try: If 10.0 g of lithium chloride decomposes how many liters of Cl2 gas are produced at STP? (write a balanced equation first!) Example 3: How many liters of H2 at 35 o C and 745 mm Hg are need to react completely with 825 g of WO3? WO3(s) + 3H2 (g) à W(s) + 3H2O (l) You try: How many L of CO at 27 0 C and 788 mm Hg can be produced from 65.5 g of carbon? 2C(s) + O2(g) à 2CO(g) Next: Practice Stoichiometry & Gas Stoichiometry Sheets

8 GAS LAW PROBLEMS Unit 11 Packet - Page 8 of 14 Work the following problems and identify the gas law used; be sure your answer includes units! 1. A gas occupies a volume of 35.9 ml at a temperature of 22.0 C. What volume will the same gas occupy at a temperature of 28.0 C? 2. At a pressure of 780 mm Hg and 24.2 C a gas has a volume of ml. What will the volume of this gas be under standard conditions? 3. A gas occupies a volume of 24.8 ml at 725 torr. What will the pressure of the gas be at 22.5 ml? 4. A gas occupies a volume of 40.8 ml at a temperature of 33.5 C. At what temperature will the volume of the gas be 39.2 ml? 5. If 45.0 ml of a gas is under 1.3 atm of pressure; at what pressure will the volume be 60.0 ml? 6. Compare the rates of effusion of nitrogen and bromine gas at the same temperature and pressure. 7. A gas at a temperature of 67.5 C and a pressure of 882 torr occupies a volume of ml. What will the volume of the gas be at 840 torr and 80.0 C? 8. A gas with a volume of ml at a pressure of atm is subjected to a pressure of atm. What is its volume at the new pressure? 9. Hydrogen gas is collected over water and the atmospheric pressure in 122kPa at 50 C. What is the partial pressure of hydrogen? Ideal Gas Law Practice Remember: PV = nrt and D = PM/RT Must use atm, L, moles and K R =.0821 L-atm/mol-K 1. How many moles of oxygen will occupy a volume of 2.5 liters at 1.2 atm and 25 0 C? 2. What volume will 2.0 moles of nitrogen occupy at 720 torr and C? 3. What pressure will be exerted by 25 g of CO2 at a temperature of 25 0 C and a volume of 500. ml? 4. At what temperature will 5.00 g of Cl2 exert a pressure of 900. torr at a volume of 750. ml? 5. What is the density of NH3 at 800. torr and C? 6. If the density of a gas is 1.2 g/l at 745 torr and C, what is its molecular mass? 7. How many moles of nitrogen gas will occupy a volume of 347 ml at 6680 torr and 27 0 C? 8. What volume will 454 grams of hydrogen occupy at 1.05 atm and 25 0 C? (remember hydrogen is diatomic!) 9. Find the number of grams of CO2 that exerts a pressure of 785 torrs at a volume 32.5 L and a temperature of 32 0 C. 10. An elemental gas has a mass of 10.3 g. If the volume is 58.4 L and the pressure is 758 torr at a temperature of C, what is the gas (Hint find the molar mass!)

9 STOICHIOMETRY WORKSHEET Unit 11 Packet - Page 9 of 14 Work the following stoichiometry problems on your own paper and attach to packet. 1. Write the correctly balanced equation for the decomposition of MgCl2. 2. Using the above equation and given 5.0 grams of MgCl2, how many grams of magnesium are produced? HINT: grams ßà moles ßà mole ratio ßà moles ßà grams 3. What is the volume of 1 mole of a gas at STP? 4. Using the equation from #1 determine the volume of chlorine gas produced from 5.0 grams of magnesium chloride. 5. Balance the following equation: Cu + HNO3 à Cu(NO3)2 + H2 6. Given 2.00 grams of copper and an excess of nitric acid how much hydrogen gas is produced? 7. How many grams of copper (II) nitrate are produced? Answer the following questions about the ideal gas law. 8. Write out the ideal gas equation. 9. What is the ideal gas constant? 10. What is the volume of a 2.00 mol gas sample at a pressure of 1.2 atm and a temperature of 301K? Gas Stoichiometry Worksheet 1. What is the density of nitrogen trioxide at STP? 2. How many moles of helium gas are there in 23 L of helium at STP? 3. What is the volume of 1.5 g of PCl gas at 545 mm Hg and 55 C? 4. How many moles of iodine gas are there in 550 ml of gas at 2.1 atm and 27 C? 5. At STP how many liters of carbon monoxide can be produced from burning 32.5 g of oxygen according to the following equation: 2 C (s) + O2 (g) à 2 CO (g) 6. Using the above equation, what volume of oxygen is required to produce 375 ml of CO at STP? 7. What volume of oxygen in liters can be collected at 750 mm Hg and 25.0 C when 36.0 g of KClO3 decomposes according to the following equation? BALANCE THE EQ FIRST! KClO3 à KCl + O2 8. What is the mass of 1258 ml of hydrogen gas at STP?

10 Mini Lab CARTESIAN DIVER Work in a group with your lab table. Each person will answer the questions after doing the lab. You do not have to do a formal lab write up. Unit 11 Packet - Page 10 of 14 Materials (per table) 1 empty plastic 2 Liter bottle tap water 1 dropper Procedure: 1. Using the prepared Cartesian Diver bottles. Squeeze the middle of the each bottle with as much force as possible. 2. Rank the bottles in the order of hardest to squeeze (requires most pressure to cause the dropper to sink) to easiest to squeeze and record your ranking order here: Questions: 1. What happened when you squeezed the bottle? Explain why this happens. (use chemistry ideas and terms) 2. Can you get the dropper to float in the middle of the bottle? How? 3. What gas law does the Cartesian diver demonstrate? Mini-lab - Marshmallow Fun! Procedure: 1. Each group will receive one syringe and two mini marshmallows (do not eat them!) 2. Pull the plunger out of the syringe. 3. Place one marshmallow in the syringe replace the plunger do not squish the marshmallow! 4. Set the syringe around 3ccs and then cover the tip of the syringe with your thumb pull back on the plunger without removing your thumb. Observe. QUESTION # 1 What happened? Use chemistry ideas in your explanation. 5. Remove your thumb reset the plunger as high as possible (around 12 ccs) put your thumb back over the syringe and press down on the plunger. QUESTION # 2 - What happened? Use chemistry ideas in your explanation. 6. Replace with a fresh marshmallow when necessary. 7. Clean up throw away used marshmallows and return syringe to your teacher. Question # 3 - What gas law does this demonstrate? Question # 4 - If the original pressure inside the syringe is 1.0 atm and the marshmallow has a volume of 1.3 ml what volume will the marshmallow occupy if the pressure is decreased to 0.75 atm? Question # 5 - Why do you eventually need a fresh marshmallow?

11 MINI LAB - THE COLLAPSING SOFT DRINK CAN Unit 11 Packet - Page 11 of 14 Do this lab as a group the whole table will work together to perform the procedure. Each student is to individually answer the questions. Procedure: 1. Place about 5 ml of water in an empty aluminum soft drink can. (Be sure the can is clean first. Place on a hot plate. 2. Boil the water in the can vigorously for a minute, until the can is filled with water vapor. (Watch for steam to escape for about 30 seconds). 3. Carefully (using tongs) and quickly invert the can and place it, top down, into a large beaker containing about 1 2 inches of water a few ice cubes. Your teacher will demonstrate how this is to be done. 4. Observe! Answer the following questions: 1. What happened when the can was placed in the water? 2. Explain why using chemistry terms! 3. What gas law does this demonstrate? Empty water from the can and place it in the can recycling. Put hot plate to the side make sure it is UNPLUGGED! Return safety glasses.

12 Unit 11 Gases, Gas Laws, & Gas Stoichiometry Unit 11 Packet - Page 12 of 14 STUDY GUIDE REVIEW INFORMATION TO BE MEMORIZED/Used: As practice; write the formulas without your notes! Charles Law Boyle s Law Gay-Lussac s Law Combined gas law Dalton s law Graham s law Ideal gas law R = Molar volume at STP STP = Density at STP and Density NOT at STP Molar Mass from the Ideal Gas Law PRACTICE PROBLEMS: Work without your notes just your reference packet! 1. If 259 ml of oxygen gas is at 112 kpa, what will the volume be at standard pressure? 2. What is the pressure of helium gas collected over water, if the barometric pressure is 88.3 kpa and the temperature is 30 C? 3. What is the volume of a gas at 273 K, if it had a volume of 22.8 ml at 48 C? 4. If a 25 ml of sulfur dioxide is at 37 C and 90.2 kpa, what will the temperature be if the volume becomes 19.6 ml and the pressure is 760 torr. 5. What is the volume of 156 grams of chlorine gas at STP? 6. What is the pressure exerted by 1.32 moles of gas in an 18 L vessel at 27 C? 7. A 759 ml vessel contains mol of a gas at 98.6 kpa. What is the temperature of the gas? 8. What volume of hydrogen gas at STP will be produced from 16.7 g of magnesium reacting with an excess amount of hydrochloric acid? 9. What volume of fluorine gas is required to react with 2.67 g of calcium bromide to form calcium fluoride and bromine at 41 C and 4.31 atm? 10. Which one of the following will diffuse the fastest at STP: NH3, CH4, Ar, HBr? ANSWERS: ml kpa ml 4. O C (273 K) L atm K L H ml 10. CH4

13 Unit 11 Packet - Page 13 of 14 Flick Your Bic Molar Mass Determination of Butane Objective: To use the ideal gas law to determine the molar mass of butane. Equipment: make an equipment list after reading the procedure Procedure: 1. Weigh the lighter. 2. Fill a graduated cylinder up with water and invert it into a large beaker/container. Make sure there are NO bubbles in the cylinder! 3. Through water displacement collect about ml of butane gas in your cylinder. You will do this by placing the lighter under the water into the mouth of the cylinder. Then press the button to release gas into the cylinder the gas should displace the water and you should get a graduated cylinder filled with gas do not light the lighter. 4. Carry your inverted cylinder with to the fume hood and release the butane in the hood do not inhale. 5. Dry off the lighter and reweigh it. The lighter must be VERY dry or the moisture will interfere with your results. 6. Use a thermometer to get the temperature of the room. 7. Check the whiteboard to get the atmospheric pressure. Mass of Butane lighter (before): Mass of Butane lighter (after): Mass of butane used: Data Table with Results (Shaded Areas are Calculated Results!) Volume of water displaced =volume of butane collected from the lighter: Temperature of the room: Atmospheric pressure: Pressure of the butane: Moles of butane: Molar mass of butane: Empirical Formula/Molecular Formula Percent error: Data Collection (don t forget units!): Calculations: Calculations: Show all work here! 1. Using Dalton s law of partial pressures calculate the partial pressure of the butane and enter in your data table beside the pressure of butane. 2. Using the ideal gas law, calculate the moles of butane and enter in your data table.

14 Unit 11 Packet - Page 14 of Using your answer for the moles of butane from (2) and your mass of butane used (data table) calculate the molar mass of butane. 4. The percentage composition of the gas in the lighter is as follows: carbon, 82.63% and hydrogen, 17.37%. Calculate the empirical formula for butane and enter in the data table. 5. Using your molar mass, see if you can get a whole-number multiple of the empirical formula for the molecular formula of the gas and enter your result in the data table. 6. Given that the molar mass of butane is 58 g/mol calculate your percent error. Questions: 1. How do you explain the fact that the gas is a liquid in the lighter, but a gas when it is collected (think phase diagram..)? Complete the following Error Analysis Table listing at least two possible errors and their effect on the molar mass (would it be higher or lower and why!) Error Analysis Conclusion: Write a paragraph conclusion a minimum of three grammatically correct sentences!