Gas Laws Notes PDF

Day 1 Chapter 4 Gases How did this happen? Gas Demo  Common Gas Terms 1) Pressure – force per unit area (standing on toes vs. flat on feet) - SI units is a kilopascal kPa (1 kPa = 1000 N/m2) - Air pressure ~ 101.3 kPa (101300 N/m2) - Other pressure unit...

Day 1 Chapter 4 Gases How did this happen? Gas Demo  Common Gas Terms 1) Pressure – force per unit area (standing on toes vs. flat on feet) - SI units is a kilopascal kPa (1 kPa = 1000 N/m2) - Air pressure ~ 101.3 kPa (101300 N/m2) - Other pressure units include (see table 2 pg. 149) a) atm – atmosphere b) mmHg – millimetre of mercury c) torr - named after Evangelista Torricelli, an Italian physicist and mathematician who discovered the principle of the barometer d) PSI – often used for tire pressure Pressure unit conversions Use a conversion factor to switch between units Example 450 Pa x 1 kPa = 0.450 kPa 1000 Pa Conversion Factor (does not change value) Complete table Pressure Conversion Factor Pressure Unit # 1 Unit # 2 6.4 atm kPa 800 torr atm 5.35 bar kPa 900 mmHg Pa 1200 Pa atm 2) Temperature Scales– a) Celsius scale – 0C (based on water’s properties) b) Kelvin scale – K (base on?) - based on absolute zero →0 K or – 273 0C Conversion → take Celsius and add 273 i.e. 25 0C → (25 + 273) → 298 K i.e. – 10 0C → (- 10 + 273) → 263 K 3) Standard Gas Conditions a) SATP – standard ambient temperature pressure 25oC and 100 kPa b) STP – standard temperature and pressure 0 oC and 1 atm (101.325 kPa) Don’t memorize, just know where to look! Chemical and Physical Properties of gases - different chemical properties (Methane vs. Nitrogen) https://www.youtube.com/watch?v=Q3stEQ3fL68 (hydrogen vs. helium balloon) - similar physical properties 1) They always fill their container (no shape or volume) 2) Highly compressible (volume depends on pressure) 3) Temperature affects pressure and the volume of a gas Inside Outside 4) They follow the gas laws H.W. 1) Graph the following Data Part One Boyle’s Law Evidence Table Pressure (kPa) Volume (L) 25 5.00 50 2.50 75 1.68 100 1.26 125 1.00 150 0.83 Plot and draw a line of best fit Part Two Charles’ Law Evidence Table Temperature (oC) Volume (L) 25 5.00 50 5.42 75 5.84 100 6.26 125 6.68 Notes: - For Part Two → make you temperature scale so it goes from – 300 0C to 150 0C - Draw a line of best fit for both and for Part two extend line until it reaches the x – axis H.W. 1) Graphs 2) Workbook pg. 2 #1-3 Day Two Looking at Graph – 1) How are Press. and Vol. related? 2) How are Temp and Vol. related?  The Gas Laws 1) Boyle’s Law (pressure/volume) - as the pressure of a gas increases, the volume of the gas decreases proportionally - Pressure and volume are inversely proportional - temp. and moles of gas (n and T) stay constant P1V1 = P2V2 (units must be the same on both sides) Example A sample of gas has a volume of 4025 mL and a pressure of 250 kPa. If the volume was expanded to 5.00L, what would the new pressure be in atmospheres? H.W. – Boyles Law 1) Workbook Pg. 3 (#4, to #8) #19 on pg. 5 2) Textbook Pg. 152 #6,7,8,9  Name:_____________________________ Boyle’s Law Investigation Problem: What effect does decreasing the pressure have on the volume of a gas? Design: M.V. R.V. Procedure: 1. Using pressure gauge, obtain pressure reading of bottle in kPa. 2. Read volume of syringe in mL. (estimate last digit) 3. Allow a small volume of gas out of bottle (0.2 mL of syringe scale). 4. Repeat steps 1-3 until approximately 10 values for pressure and volume have been recorded. 5. Obtain the value for atmospheric pressure and add that value to the pressures measured to get the adjusted pressure. 6. Multiply the adj. Pressure by the volume for each set of data to get a constant. Use only whole number values for k. Evidence Trial Pressure Volume Adj. Press. PxV = k 1 2 3 4 5 6 7 8 9 10 Analysis: 1. Based on your experimental evidence, what is the answer to the lab’s problem? 2. graph the volume vs. the adj. Pressure. Compare to the graph from day 1 notes. (graph paper on reverse side of page) Evaluation: 1. Is Boyle’s law verified? Calculate the % difference between the two most widely separated values for “k” to determine verification. 2. Are there any sources of error, possible improvements? Day 3 Remember: How are Vol. and Temp. related? (graph) 2) Charles’ Law (volume / temperature) - The volume of a gas increases proportionally to the absolute temperature V 1 = V2 ( Units must be the same on both sides Temp. in Kelvins) T 1 T2 - pressure and the moles of gas (n and P) stay constant Formula page – T1 = i.e. A sample of gas at 25 0C occupies a volume of 4.0 L. The temperature is increased 35 0C. What is the new volume of gas? H.W. 1) Workbook #9, to #12 2) Textbook Pg. 156 #14,15,16,17 3) The Combined Gas Law - the amount or moles of gas “n” is constant P1V1 = P2V2 (temperature must be in Kelvins) T1 T2 A balloon contains hydrogen gas at 20 0C and 100 kPa and has a volume of 7.50 L. Calculate the volume of the balloon if the temperature changes to – 36 0C and the pressure changes to 28 kPa. Could there be another gas law? P1 = P2 T1 T2 Volume and amount must remain constant H.W. 1) Workbook – #1 - #18 and #19 (formula practice) 2) Pg. 159 #19,20,21,22,23 Day 4 Kinetic Molecular Theory pg. 163 1) Gases move randomly and continually. Molecules collide with each other and the walls of the container…perfectly elastic collisions. What is pressure? - based on collisions with container walls a) force of collisions b) area of collisions are taking place over Circle larger Pressure for each? or cold warm big area small area (5 molecules at 10 0C vs 50 0C) (5 molecules at same temp) 2) Gas molecules are far apart from each other 20 - 30 times the size of molecule. Why? Weak intermolecular forces. Compressible 3) Attractive forces between gas molecules are weak and usually negligible (except at low temp. and high pressure) If you remove lid from container, what happens?  4) The average kinetic energy of the molecules in a gas depends on temperature. 1) Which box is cold / warm? 2) Which box is high / low pressure? H.W. 1) Workbook #20 - #31 (pg. 6) 2) Describe - Using KMT - Charles’ and Boyle’s Law - liquid water to gas (phase change) Boyle’s Law and Kinetic Molecular Theory (KMT) The kinetic theory explains Boyle's law Why does the pressure of a gas increase when the volume of the container decreases? Remember that the pressure of a gas on the walls of the container is due to the collisions of the molecules on the walls of the container. The change in momentum of these molecules in unit time is a force exerted by the walls of the vessel on the molecules, which, by Newton's third law, exert an equal and opposite force on the walls of the vessel. This force, divided by the area of the walls in contact with the gas, is the pressure of the gas. If the volume of the container is reduced, the gas molecules have a shorter distance to travel before they collide with the walls of the container. This means that they collide with the walls more frequently. The change of momentum that takes place at the surface of the walls occurs in a shorter time, resulting in a greater force exerted by the molecules on the walls, and hence a greater pressure.  Law of Combining Volumes pg. 164 - a method to determine the volume of gaseous reactants or products - Pressure and temp must remain constant Steps 1) write balanced chemical equation 2) use V2 = V1 n2 n = moles n1 P1 V1 = P2 V2 (Press. and Temp are constant) T1 n1 T2 n2 Thus, V1 = V2 (always solve for V2 ) n1 n2 Example – Production of ammonia gas from nitrogen and hydrogen at SATP. If 1.6 L of N2 (g), what is volume of other gases 1 N2 (g) + 3 H2 (g) → 2 NH3 (g) 1.6 L ? ? Example Predict the volume of gaseous products if 180 mL of oxygen is used when ethane is combusted. Examples Calculate the volume of reactants and products if 3.0 L of propane is burned H.W. 1) Workbook Pg. 10 # 51 - 60  Day 5 Avogadro’s Theory pg.165 Gases are far apart from each other (20-30 times their size) - thus, equal volumes of gases (eg. N2 or H2 or NH3) at same temp. and pressure contain equal number of molecules or 2.0 L = 1 mole H2 2.0 L = 1 mole NH3 Temp. = 150 K and Press. = 50 kPa All balloons at SATP conditions Volume to mole Conversion (using molar volume) Given Unit X Conversion Factor = Required Unit Conversion Factor: - Equal to 1 - Will be a fraction - Volume to Moles conversion us Molar Volume What is molar volume? 1 mole of any gas at STP = 22.7 L 1 mole of any gas at SATP = 24.8 L V = molar volume V @ SATP = 24.8 L/mol V @ STP = 22.7 L/mol Steps to Solving a molar volume question a) is a mass to mole conversion required? b) Conversion factors → V (VSATP or VSTP ) c) do required calc. Note – this is not a before and after situation Examples How many moles of oxygen are there in a volume of 5.6 L at STP? How many mL of methane at SATP is 2.15 moles?  The Mole Map We can link mass/mole conversions to molar volume Example What is the volume occupied by 3.50 g of helium @ SATP? H.W. 1) Workbook pg. 7 #32 - 35 2) Pg. 171 # 3, 5 to 9 Day 6 1.Complete The Chart Law Relationship Graph Boyle’s Law P1V1 = P2V2 Charle’s Law Gay-Lussac’s Law Combined Gas Law Law of Combining Volume 2. Describe Boyle’s, Charles’ and Lussac’s Law using KMT. 3. Butane gas is combusted by accident in tomorrow’s lab. Calculate the volume of water vapour produced if 825 mL of oxygen is used. 4. A kid gets a balloon for their birthday. The room is around SATP conditions. As the kid walks to the car (outside temp. around -25 0C), something happens to his balloon. a) Describe what is happening to a gas molecule in side the balloon using KMT b) Verify this using the appropriate gas law calculation 5. Why does a bag of chip expand as a hiker walks up a mountain? Assume temperature is constant. 6. On a cold morning I noticed this appeared on my dashboard. It was not there the previous day. Explain using a gas law. (volume assumption of a tire?) Boyles Law P1V1 = K (moles + temp. don’t change) Charles Law V1 = K (moles + Press. don’t change) T1 Combined Gas Law P1V1 = K (moles don’t change) T1 Law of Combining Volumes V1 = V2 (Temp. + Press. don’t change) n1 n2  The Ideal Gas Law P1V1 = Constant “R” (Ideal Gas Constant) n1T1 Usually written as pv = nRT Formula pv = nRT p = pressure kPa v = volume L n = moles R = universal gas constant T = temp. K Calculate the mass of neon that should be added to a tube that is 1.2 L big to produce a pressure of 120 kPa at 40 0C. (1.1 g) H.W. 1) Workbook pg. 9 # 45 – 50 2) pg. 176 # 2, 5, 6, 7, 9, 3) Pre Lab Inv. 4.3 What is an ideal gas? - ideal gas – hypothetical gas that obeys all gas laws perfectly under all conditions - does not condense into liquids (high press. or low temp.) Ideal Student? vs. A Real student? - real gases deviate most at high pressure and low temperatures (KMT?) Summary Ideal Gas Real Gas - particle has no attraction - do the opposite of - do not condense as press. ideal gas increase and temp. decreases - all collisions are perfectly elastic In chem. 20 – we will deal with gases as if they were ideal H.W. Pg. 8 #36 - 44  Inv. 4.3 Using The Ideal Gas Law Pg. 179 Problem: What is the molar mass of Butane gas? Prediction: Procedure: Evidence Table: Molar Mass of Butane Collection of Data Variable Value Mass initial Mass final Mass of gas Pressure Vol. of gas Temperature of water Analysis: 1. Use the ideal gas law to determine the molar mass of butane gas experimentally. 2. Determine the percent difference between your experimental result and the expected result. 3. Why does this lab design work for butane gas? Evaluation: 1. Evaluate your prediction. 2. Sources of error. 3. Effects of sources of error. Day 7 Inv. 4.3 Using The Ideal Gas Law Demo pg. 179 Problem: copy Design: see page. 179 Prediction: Calculate the molar mass of butane C4H10 (g) Procedure: Read and summarize Evidence: create a table to record results from lab (see procedure) Analysis: 1.Use the ideal gas law to determine the molar mass of butane. 2.Determine the percent difference between your experimental result and the expected result. Evaluation: 1. Evaluate your prediction. (10% guideline) 2. Sources of error. 3. Effects of sources of error.  Butane Lab Marking Guide Problem (1) C N - copied from text book Prediction (1) C N - molar mass of butane with correct units Procedure (1) C P N - numbered steps - safety is mentioned (goggles and butane disposal) - easy to follow - cleanup is mentioned Evidence (1) C P N - table format to record results - ruler used, title - values fit into cells - a complete table Analysis (4) C P N 1) Molar mass of Butane - formula rearranged - values substituted with units - ans. with correct units 2) Percent error calculation Evaluation (1) C P N - evaluation of prediction - sources of error - effect of sources of error Total / 9 Day 8 Density Would Argon balloons work at a birthday party? Why are balloons filled with Helium? Why does a CO2 fire extinguisher work to smother a fire? Why does a hot air balloon rise? Phase Changes and KMT How could you describe the water molecules in each phase?  Gas Stoichiometry How it’s done: we will follow 6 steps: 1) write a balanced equation 2) determine given and wanted 3) mole conversion (moles of given) 4) mole Ratio 5) moles to ? conversion (exit the question) What volume of ammonia at 400 kPa and 80 0C can be obtained from the complete reaction of 7.5 L of hydrogen @ STP? If 300 g of propane burn in a gas barbecue, what volume of oxygen at SATP is required for the reaction? H.W. #61 - #72