Around 1827, dust particles were seen to move in a random, zig-zag pattern under a microscope,

1. Chapter 13“GASES” Brownian Movement Around 1827, dust particles were seen to move in a random, zig-zag pattern under a microscope, “I’ve been behind this guy in the hall!”

2. From the idea of Brownian Movement came the explanation for the behavior of gases and, later, for other particles of matter. So, let’s look at some important properties of gases, shall we? Why sure….

3. 1. Gases have mass. • Gases seem to be weightless, but they are classified as matter, which means they have mass. • The density of a gas – the mass per unit of volume – is much less than the density of a liquid or solid, however.

4. Lower density Higher density GAS DENSITY 22.4 L of ANY gas AT STP = 1 mole

5. 2nd– Gases are Compressible • If you squeeze a gas, its volume can be reduced considerably • The low density of a gas means there is a lot of empty space between gas molecules.

6. 3rd – Gases fill their containers • Gases spread out to fill containers until the concentration of gases is uniform throughout the entire space. • This is why there is never an absence of air around you!

7. 4th – Gases diffuse Because of all of the empty space between gas molecules, another gas molecule can pass between them until each gas is spread out evenly throughout the entire container. This is called diffusion.

9. 5th – Gases exert pressure • Gas particles exert pressure by colliding with objects in their path. • The sum of all of the collisions makes up the pressure the gas exerts.

10. Imagine a gas in a container as a room of hard rubber balls. • The collisions of the balls bouncing around exert a force on the object that with which they collide. • The definition of Pressure is force per unit area – so the total of all of the tiny collisions makes up the pressure exerted by the gas.

11. 6th – Pressure depends on Temp • The higher the temperature of a gas -the higher the pressure that the gas exerts • The reverse of that is true as well, a the temperature of a gas decreases – the pressure decreases. • How are temperature & pressure related? • DIRECTLY

12. Kinetic Molecular Theory: Gases • particles in continuous, random, rapid motion (Brownian!) • collisions between particles are elastic (they may transfer energy between them, but they don’t lose any energy) • volume of the particles is negligible (little effect on their behavior) • No attractive forces between particles (littleeffect on their behavior)

13. What kind of energy is “energy of motion”? That’s Right! Kinetic Energy! What can you say about particles that have a lot of kinetic energy? Right again! They move Fast! But the particles in a gas don’t all move at the same speed so we measure theAverage Kinetic Energy of a sample – TEMPERATURE!

14. Let’s think of your average, room temperature oxygen molecule as a car! It would be traveling at a speed of about 443 m/sec (1000 mil/hr)! If we compare speed & collisions, our “Oxygen Car” would have a collision every 314 car lengths (or 4.5 x 109 collisions per second)! Tough getting insurance!

15. Gas variables • In order to describe a gas sample completely and then make predictions about its behavior under changed conditions, it is important to deal with the values of: 1) pressure 2) temperature 3) volume 4) amount of gas

16. Pressure (P) • The pressure of a gas is the force exerted on the wall of the container a gas is trapped in. • It is the force of the collisions & the number of collisions with the walls of a container that cause Gas Pressure. • No Gas – No Pressure – VACUUM!

17. Put on your bathing suits (NO SPEEDOS, Please!) and I’ll meet you on a sandy beach in the Bahamas! 1, 2, 3… Let’s go! You Look Great! Now swim out to that buoy!

18. When I say go, take a deep breath & dive under water. Go down about 2 feet & look up. Then come back to the surface! What did you see directly above you? Water! Was it heavy? Did you even notice?

19. Now, take a deeper breath & go down About 10 feet and look up. Then come up quickly! What did you notice above you this time? Yes! More water! Did you feel it this time? How about those ears?!

20. Swim to shore! We’ll take the rest of our notes there! We live under a sea of air. The more air above you, the more pressure on you! There is less air above a mountain top than There is above a valley, so, less pressure above the mountain!

21. Even though babies know how to drink from a straw, most people, young or old, don't know how it works. Most people think the suction caused by our mouth pulls the liquid up through the straw.

22. Normally when your mouth is closed, there's not much air (blue spheres) inside your mouth. They bounce around causing 15 psi pressure in the mouth. However, when you drop your jaw and keep your lips closed, there's more room for the air to spread out . so the air pressure inside the mouth is less. Outside air, at a higher pressure, is trying to get into the mouth. It pushes on the cheeks causing them to be sunken.

23. If you have a straw in your mouth, then air pressure pushing on your drink has more force than the force from the air in your mouth. The outside air pressure pushes onto the surface of the drink. This pressure pushes liquid up through the straw to your mouth.

24. Normal Air (atmospheric) Pressure is the average pressure of the air at sea level under normal conditions. It will support a column of mercury 760 mm high

25. Air pressure is measured with a Barometer There are several units for pressure depending on the instrument used to measure it including: 1) atmospheres (atm) 2) Millimeters of Mercury (mmHg) 3) Kilopascal (kPa) Remember STP? 0° C & 1 atm 1atm = 760 mm Hg = 101.3 kPa

26. Pressure Conversions A. What is 475 mm Hg expressed in atm? 1 atm 760 mm Hg B. The pressure of a tire is measured as 233 kPa. What is this pressure in mm Hg? 760 mm Hg 101.3 kPa 475 mm Hg x = 0.625 atm 233 kPa x = 1.75 x 103 mm Hg

27. Temperature (T) • The temperature of a gas is generally measured with a thermometer in Celsius. • All calculations involving gases should be made after converting the Celsius to Kelvin temperature. Kelvin = C° + 273

28. Volume (V) • The volume of the gas is simply the volume of the container it is contained in. • The metric unit of volume is the liter (L) 1 L = 1 dm3 = 1000 ml = 1000 cm3

29. Amount (n) • The quantity of gas in a given sample is expressed in terms of moles of gas (n). • This of course is in terms of 6.02 x 1023 molecules of the gas. • Don’t forget to convert mass to moles you just divide by the molar mass of the gas.

30. Gas Laws • Studies of the behavior of gases played a major role in the development of physical sciences in the 17th and 18th centuries. • The Kinetic Molecular theory marked a significant achievement in understanding the behavior of gases. • Observations have become mathematical laws which we can use to predict quantitative outcomes.

31. (twice as many molecules) Avogadro’sHypothesis Equal volumes of gases at the same T and P have the same number of molecules. V and n are directly related. 2 moles 2 x avogadro’s # 44.8 L at STP 1 mole of gas 6.02 x 1023 atoms 22.4 L at STP

32. Avogadro’s Hypothesis and Kinetic Molecular Theory The gases in this experiment are all measured at the same T and V. P is also directly proportional to number of particles (n).

33. Dalton’s Law John Dalton 1766-1844

34. Dalton’s Law of Partial Pressures • At constant temperature, the pressure of a mixture of gases that do not react equals the sum of the partial pressures of the gases in the mixture. • Ptotal(in gas mixture) = P1 + P2 + P3...

35. Dalton’s Law of Partial Pressure

36. Dalton’s Law of Partial Pressures • 2 H2O2 (l) ---> 2 H2O (g) + O2 (g) • 0.32 atm 0.16 atm What is the total pressure in the flask? Ptotal = PH2O + PO2 = 0.48 atm

37. Collecting a gas “over water”

38. Collecting Gases over Water Total gas pressure = 769 mm Hg Water vapor Pressure @ 19C= 17 mmHg Total pressure – Water vapor pressure = Pressure of gas 769 – 17 = 752 mm Hg

39. Boyle’s Law • Robert Boyle was among the first to note the relationship between pressure and volume of a gas. • He measured the volume of air at different pressures, and observed a pattern of behavior which led to his mathematical law • During his experiments temperature and amount of gas weren’t allowed to change Robert Boyle (1627-1691). Son of Earl of Porrello, Ireland.

40. Boyle’s Law The volume of a confined gas varies inversely with the pressure, if temperature remains constant. P goes up as V goes down. P1V1 = P2 V2

41. As the pressure increases Volume decreases

42. Boyle’s Law and Kinetic Molecular Theory P proportional to 1/V

43. Boyle’s Law: What if we had a change in conditions? since PV = k (k = constant) P1V1 = P2V2 Ex: A gas has a volume of 3.0 L at 2 atm. What is its volume at 4 atm?

44. determine which variables you have: • P1 = 2 atm • V1 = 3.0 L • P2 = 4 atm • V2 = ? • determine which law is being represented: P and V = Boyle’s Law

45. P1V1 = V2 P2 (2.0 atm)(3.0L) = V2 (4atm) 3) Rearrange the equation for the variable you don’t know 4) Plug in the variables and chug it on a calculator: V2 = 1.5L

46. Volume Pressure How does Pressure and Volume of gases relate graphically? (typical graph for an inverse proportion) PV = k (Temperature, # of particles remain constant)

47. Charles’s Law • Jacques Charles determined the relationship between temperature and volume of a gas. • He measured the volume of air at different temperatures, and observed a pattern of behavior which led to his mathematical law. • During his experiments pressure of the system and amount of gas were held constant.

48. Charles’s Law The volume of a gas varies directly with the absolute temperature, if pressure remains constant. Pressure & Kelvin Temp. are directly proportional! V1 V2 Jacques Charles (1746-1823). Isolated boron and studied gases. Balloonist. = T1 T2

49. Volume of balloon at room temperature Volume of balloon at 5°C

50. V1 V2 = T1 T2 Charles’s Law: What if we had a change in conditions? since V/T = k Eg: A gas has a volume of 3.0 L at 127°C. What is its volume at 227 °C?