Chapter 14

1. Chapter 14 Gases And Gas Laws

2. Gas Behavior • Certain gas behaviors are well known: 1. Gases can be compressed by exerting pressure on them 2. Gases will occupy the space they are in.

3. There are certain Physical Properties that all gases share. 1. Gases have mass. -empty vs. filled balloon 2. It is easy to compress a gas. • If you squeeze a gas, its volume can be reduced considerably. • Shock absorber

4. 3. Gases fill their containers completely • When a balloon is filled with air, the air is distributed evenly throughout the balloon. • At room temperature, the distance between particles is about 10x the diameter of the particle.

5. 3. cont. • How does the volume of the particles in a gas compare to the overall volume of the gas?

6. 4. Different gases can move through each other quite rapidly through diffusion. • You are observing gas diffusion when you smell food cooking while you are still in your bedroom.

7. 5. Gases exert pressure. -Airplane……..ears popping -blowing up a balloon 6. The pressure of a gas depends on its temperature. -Higher the temperature, more pressure because the particles are moving faster and hitting the sides of the container at a harder rate. -Lower temperature, less pressure

8. Gas properties are explained by the Kinetic Molecular model that describes the behavior of the submicroscopic particles that make up a gas.

9. Kinetic Molecular Theory 1. A gas consists of very small particles, each of which has a mass. 2. The distances separating gas particles are relatively large. 3. Gas particles are in constant, rapid, random motion.

10. 4. Collisions of gas particles with each other or with the walls are perfectly elastic. -Elastic collisions: never loses momentum (bouncing ball example). 5. Gases have a high Kinetic energy at higher temperatures because the molecules are moving faster. They also have a lower Kinetic Energy at lower temperatures.

11. 6. Gases DO NOT exert force on one another.

12. Measuring Gases • Variable Symbol Units • Amount of a gas n mol • Volume V L, mL 1000mL = 1L 3. Temperature T K, Kelvin (celcius + 273) 4. Pressure P atm, mmHg, KPal

13. *STP = Standard Temperature and Pressure = O degrees celcius and 1 atm. • *1 atm = 760 mmHg = 101, 325 KPa

14. I. Pressure and the number of molecules are directly related • More molecules means more collisions. • Fewer molecules means fewer collisions. • Gases naturally move from areas of high pressure to low pressure because there is empty space to move into – a spray can is example.

15. Common use? • Aerosol (spray) cans • gas moves from higher pressure to lower pressure • a propellant forces the product out • whipped cream, hair spray, paint • Is the can really ever “empty”?

16. II. Volume of Gas • In a smaller container, the molecules have less room to move. • The particles hit the sides of the container more often. • As volume decreases, pressure increases. (think of a syringe)

17. III. Temperature of Gas • Raising the temperature of a gas increases the pressure, if the volume is held constant. • The molecules hit the walls harder, and more frequently! • Should you throw an aerosol can into a fire? What could happen? • When should your automobile tire pressure be checked?

18. The Gas Laws • These will describe HOW gases behave. • Gas behavior can be predicted by the theory. • The amount of change can be calculated with mathematical equations. • You need to know both of these: the theory, and the math

19. Robert Boyle(1627-1691) • Boyle was born into an aristocratic Irish family • Became interested in medicine and the new science of Galileo and studied chemistry.  • A founder and an influential fellow of the Royal Society of London • Wrote extensively on science, philosophy, and theology.

20. #1. Boyle’s Law - 1662 Gas pressure is inversely proportional to the volume, when temperature is held constant. • Pressure x Volume = a constant • Equation: P1V1 = P2V2 (T = constant)

21. Graph of Boyle’s Law

22. Jacques Charles (1746-1823) • French Physicist • Part of a scientific balloon flight on Dec. 1, 1783 – was one of three passengers in the second balloon ascension that carried humans • This is how his interest in gases started • It was a hydrogen filled balloon – good thing they were careful!

23. #2. Charles’s Law - 1787 • The volume of a fixed mass of gas is directly proportional to the Kelvin temperature, when pressure is held constant. • This extrapolates to zero volume at a temperature of zero Kelvin.

24. Converting Celsius to Kelvin • Gas law problems involving temperature will always require that the temperature be in Kelvin. (Remember that no degree sign is shown with the kelvin scale.) • Reason? There will never be a zero volume, since we have never reached absolute zero. Kelvin = C + 273 °C = Kelvin - 273 and

25. Joseph Louis Gay-Lussac (1778 – 1850) • French chemist and physicist • Known for his studies on the physical properties of gases. • In 1804 he made balloon ascensions to study magnetic forces and to observe the composition and temperature of the air at different altitudes.

26. #3. Gay Lussac’s Law - 1802 • The pressure and Kelvin temperature of a gas are directly proportional, provided that the volume remains constant. • How does a pressure cooker affect the time needed to cook food?

27. #4. The Combined Gas Law The combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas.

28. The combined gas law contains all the other gas laws! • If the temperature remains constant... P1 V1 P2 x V2 x = T1 T2 Boyle’s Law

29. The combined gas law contains all the other gas laws! • If the pressure remains constant... P1 V1 P2 x V2 x = T1 T2 Charles’s Law

30. The combined gas law contains all the other gas laws! • If the volume remains constant... P1 V1 P2 x V2 x = T1 T2 Gay-Lussac’s Law

31. #5. The Ideal Gas Law • Equation: P x V = n x R x T • Pressure times Volume equals the number of moles (n) times the Ideal Gas Constant (R) times the temperature in Kelvin. • R = 8.31 (L x kPa) / (mol x K) • R = 0.0821 (L x atm)/(molxK)

32. Ideal Gases • We are going to assume the gases behave “ideally”- in other words, they obey the Gas Laws under all conditions of temperature and pressure • An ideal gas does not really exist, but it makes the math easier and is a close approximation. • Particles have no volume? Wrong! • No attractive forces? Wrong!

33. Ideal Gases • There are no gases for which this is true; however, • Real gases behave this way at a) high temperature, and b) low pressure. • Because at these conditions, a gas will stay a gas!

34. Real Gases and Ideal Gases

35. Ideal Gases don’t exist, because: • Molecules do take up space • There are attractive forces between particles - otherwise there would be no liquids formed

36. Real Gases behave like Ideal Gases... • When the molecules are far apart. • The molecules do not take up as big a percentage of the space • We can ignore the particle volume. • This is at low pressure

37. Real Gases behave like Ideal Gases… • When molecules are moving fast • This is at high temperature • Collisions are harder and faster. • Molecules are not next to each other very long. • Attractive forces can’t play a role.

38. Diffusion is: • Molecules moving from areas of high concentration to low concentration. • Example: perfume molecules spreading across the room. • Effusion: Gas escaping through a tiny hole in a container. • Both of these depend on the molar mass of the particle, which determines the speed.

39. Diffusion:describes the mixing of gases. The rate of diffusion is the rate of gas mixing. • Molecules move from areas of high concentration to low concentration.

40. Effusion: a gas escapes through a tiny hole in its container -Think of a nail in your car tire… Diffusion and effusion are explained by the gas laws.

41. Types of Gas Laws 1. Boyle’s Law • P1V1 = P2V2 • Pressure and Volume have an inverse relationship. • If one increases, then the other decreases.

42. 2. Charles’s Law • V1 = V2 • T1 T2 • Volume and temperature have a direct relationship. • Both are either increasing or decreasing at the same time.

43. 3. Ideal Gas Law • PV = nRT • R is a constant

44. Gay-Lussac’s Law • P1 = P2 • T1 T2

45. End of Chapter 13