Gas Laws

1. Gas Laws Charles, Boyle, Gay-Lussac, Combined and The Ideal Gas Law

2. What did you learn from the lab? • What we know about gases from observation • Gases have mass. • Its easy to compress gases. • Gases fill their containers completely. • Different gases can move through each other quite rapidly. • Gases exert pressure. • The pressure of a gas depends on its temperature.

3. Kinetic-Molecular Theory The KINETIC MOLECULAR THEORY is based on the idea that particles of matter are always in motion. It says… • Gases have no volume • Gases are in constant, random motion. • Gases to not attract or repel each other. • Gases have elastic collisions. • The average kinetic energy of gas particles is directly related to their temperature.

4. Variables That Effect Gases • Moles (n) – the amount of gas. • Volume (V) – the size of the container that holds the gas in liters (L). • Temperature (T) – the speed or kinetic energy of the particles in kelvin (oC +273) • Pressure (P) – The outward push of gas particles on their container in atmospheres (atm), millimeters of mercury (mm Hg), kilopascals (kPa) or torr (torr). *Think of pressure as the number and strength of collisions between gas particles and their container.

5. FACTORS THAT AFFECT THE BEHAVIOR OF GASES • PRESSURE (P) - the dance movement • TEMPERATURE (T) - beat of the music • VOLUME (V) - size of room • AMOUNT (n) - # of dancers

6. STP • The behavior of a gas depends on its temperature and the pressure at which the gas is held. • So far we have only dealt with gases at STP. Standard Temperature and Pressure. • 273 kelvins and 1 atm.

7. UNIT HELP! – Pressure Units • 1 standard atmosphere (atm) = 760 mm Hg (torr) = 29.2 inches Hg = 101.3 kPa

8. UNIT HELP! • Temp Units • Celsius • Kelvin = C + 273.15 • Volume Units • mL, L, cm3 • n Units • moles

9. The Gas Laws • Boyle’s Law • Charles’s Law • Gay-Lussac’s Law • The Combined Gas Law • The Ideal Gas Law

10. Boyle’s Law • The Pressure-Volume Relationship • The pressure of a fixed mass of gas varies indirectly with the volume • What remains constant? • What changes? (As one goes up, the other goes down) • P1V1 =P2V2 • If 3 of the variables are known, the fourth can be calculated.

11. BOYLE’S LAW

13. Boyle’s Law • The gas in a 20.0mL container has a pressure of 2.77atm. When the gas is transferred to a 34.0mL container at the same temperature, what is the new pressure of the gas. P1V1 =P2V2

14. Boyle’s Law • So, does it make sense? • If a set amount of gas is transferred into a larger container, would the pressure go up or down? • Would there be more collisions, or fewer collisions with the container holding the gas? • More volume (space) means fewer collisions with the container, therefore pressure goes down. (From 2.77 atm to 1.63 atm)

15. Charles’s Law • The temperature-volume relationship • The volume of a fixed mass of gas varies directly with the temperature. • What remains constant? • What changes? (As one goes up, the other goes up.) If 3 of the variables are known, the fourth can be calculated.

16. CHARLES’ LAW ANIMATION ANIMATION#2

18. Charles’s Law • What will be the volume of a gas sample at 355K if its volume at 273K is 8.57L?

19. Charles’s Law • Does it make sense? • If the temperature of a given quantity of gas is increased, what will happen to the volume it occupies? (In an elastic container?) • Gas particles moving faster would have more collisions with the container and exert more force to enlarge the volume of the elastic container. • In this case, from 8.57L to 11.1L.

20. Gay-Lussac’s Law • The Temperature-Pressure Relationship • The pressure of a fixed mass of gas varies directly with the temperature at a constant volume. • What is constant? • What is changing? (If one goes up, the other goes up.) • If you know 3 of the variables, you can calculate the 4th.

21. Gay-Lussac’s Law • The gas left in a used aerosol can is at a pressure of 2.03atm at 25oC. If this can is thrown onto a fire, what is the pressure of the gas when its temperature reaches 928oC?

22. Gay-Lussac’s Law • Does it make sense? • If the temperature of a fixed amount of gas goes up, the particles will have more collisions. More collisions means the pressure will increase. • In this case, when the temp went up the pressure increased from 2.03atm to 8.18atm.

23. PT V P T V PT V 1 V Boyle’s P a ___ a Charles V T a Gay-Lussac’s P T Pressure - Temperature - Volume Relationship

24. The Combined Gas Law • Expresses the relationship between pressure, temperature, and volume of a fixed amount of gas. • What is constant? • 5 of the variables must be known to calculate the 6th.

25. The Combined Gas Law • The volume of a gas-filled balloon is 30.0L at 40oC and 1.75atm of pressure. What volume will the balloon have at standard temperature and pressure?

26. The Combined Gas Law • Does it make sense? • You have a fixed volume of gas. The temperature decreases which would cause fewer collisions and the pressure decreases which causes fewer collisions as well. What can you do to volume to make the pressure decrease??? • Increase it. More space means fewer collisions.

27. The Ideal Gas Law • Describes the physical behavior of an ideal gas in terms of the pressure, volume, temperature and the number of moles of gas. • Ideal – a gas as it is described by the kinetic-molecular theory postulates. • All gases are REAL gases… which behave like ideal gases only under most ordinary conditions.

28. The Ideal Gas Law • Only at very low temperatures and very high pressures do real gases show significant non-ideal behavior. • We will assume that gases are close to ideal and that the ideal gas equation applies.

29. Ideal Gas Equation PV=nRT • P-pressure • V-volume • n-number of moles of gas • R-ideal gas constant (universal gas constant) 0.0821 atm.L/mol.K 8.314 kPa . L/mol.K 62.396 torr.L/mol.K 62.4 mmHg.L/mol.K • T-temperature

30. Ideal Gas Equation • What is the volume occupied by 9.45g of C2H2 at STP? First, calculate amount of gas in moles.

31. Ideal Gas Law

32. Ideal Gas Law • How many moles of a gas at 100oC does it take to fill a 1.00L flask to a pressure of 1.5atm?

33. Lifting Power of Gases • For a gas to be used to inflate lighter-than-air craft like balloons and blimps, the gas must have a density lower than air. • The lower the density, the greater the lifting power.

34. Lifting Power of Gases • The density of a gas depends on its pressure, temperature and molar mass. • Each of these variables is part of the ideal gas law. • Therefore, we should be able to adjust each of these variables to give low density.

35. Lifting Power of Gases • However, if the pressure of the gas within a balloon or blimp were significantly less than the atmospheric pressure, the balloon or blimp would be crushed. • Therefore, only two factors can be manipulated to lower the density of a gas: molar mass and temperature

36. Molar Mass • Gases with low density can be corrosive, combustible, flammable or chemically active in some way. These gases would make poor choices to fill blimps and balloons. • Helium, due to its small molar mass and chemical inactivity is the primary choice for balloons and blimps.

37. Temperature • Helium is relatively rare and very expensive, so hot air is often preferable. • As the temperature of a gas is increased, the particles increase the number of collisions and increase the pressure inside the balloon. The volume of the balloon increases and becomes less dense and rises. • Hot air does not have the same lifting power as helium, but it is much cheaper.

38. Gas Effusion • The movement of atoms or molecules through a hole so tiny that they do not stream through but instead pass through one particle at a time. • Explains why helium balloons deflate slowly over a period of a few hours. • The lower the mass of the gas, the greater the speed of its particles. • Hydrogen effuses faster than helium. Helium effuses faster than oxygen.