Chemistry - Chp 14 - The Behavior of Gases - PowerPoint


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Chemistry - Chp 14 - The Behavior of Gases - PowerPoint

  1. 1. Chapter The 14 Behavior of Gases
  2. 2. Section 14.1 The Properties of Gases OBJECTIVES: Explain why gases are easier to compress than solids or liquids are.
  3. 3. Section 14.1 The Properties of Gases OBJECTIVES: Describe the three factors that affect gas pressure.
  4. 4. Compressibility Gases can expand to fill its container, unlike solids or liquids The reverse is also true:  They are easily compressed, or squeezed into a smaller volume  Compressibility is a measure of how much the volume of matter decreases under pressure
  5. 5. Compressibility Thisis the idea behind placing “air bags” in automobiles  In an accident, the air compresses more than the steering wheel or dash when you strike it  The impact forces the gas particles closer together, because there is a lot of empty space between them
  6. 6. Compressibility Atroom temperature, the distance between particles is about 10x the diameter of the particle  Fig. 14.2, page 414 This empty space makes gases good insulators (example: windows, coats) How does the volume of the particles in a gas compare to the overall volume of the gas?
  7. 7. Variables that describe a Gas The four variables and their common units: 1. pressure (P) in kilopascals 2. volume (V) in Liters 3. temperature (T) in Kelvin 4. amount (n) in moles• The amount of gas, volume, and temperature are factors that affect gas pressure.
  8. 8. 1. Amount of Gas When we inflate a balloon, we are adding gas molecules. Increasing the number of gas particles increases the number of collisions thus, the pressure increases If temperature is constant, then doubling the number of particles doubles the pressure
  9. 9. Pressure and the number of molecules are directly related More molecules means more collisions, and… 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.
  10. 10. Common use?A practical application is Aerosol (spray) cans  gas moves from higher pressure to lower pressure  a propellant forces the product out  whipped cream, hair spray, paint Fig. 14.5, page 416 Is the can really ever “empty”?
  11. 11. 2. 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)  Thus,volume and pressure are inversely related to each other
  12. 12. 3. Temperature of Gas Raising the temperature of a gas increases the pressure, if the volume is held constant. (Temp. and Pres. are directly related)  The molecules hit the walls harder, and more frequently! Fig. 14.7, page 417 Should you throw an aerosol can into a fire? What could happen? When should your automobile tire pressure be checked?
  13. 13. Section 14.2 The Gas Laws OBJECTIVES: Describe the relationships among the temperature, pressure, and volume of a gas.
  14. 14. Section 14.2 The Gas Laws OBJECTIVES: Use the combined gas law to solve problems.
  15. 15. The Gas Laws are mathematical Thegas laws 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
  16. 16. Robert Boyle • Boyle was born into an aristocratic Irish family(1627-1691) • 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.
  17. 17. #1. Boyle’s Law - 1662Gas pressure is inversely proportional to thevolume, when temperature is held constant. Pressure x Volume = a constant Equation: P1V1 = P2V2 (T = constant)
  18. 18. Graph of Boyle’s Law – page 418 Boyle’s Law says the pressure is inverse to the volume. Note that when the volume goes up, the pressure goes down
  19. 19. - Page 419
  20. 20. 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!
  21. 21. #2. Charles’s Law - 1787The volume of a fixed mass of gas isdirectly proportional to the Kelvintemperature, when pressure is heldconstant.This extrapolates to zero volume at atemperature of zero Kelvin. V1 V2 = ( P = constant) T1 T2
  22. 22. Converting Celsius to Kelvin•Gas law problems involvingtemperature will always require thatthe temperature be in Kelvin.(Remember that no degree sign isshown with the kelvin scale.) •Reason? There will never be a zero volume, since we have never reached absolute zero.Kelvin = °C + 273 and °C = Kelvin - 273
  23. 23. - Page 421
  24. 24. Joseph Louis Gay-Lussac (1778 – 1850) French chemist andphysicist Known for his studies onthe physical properties ofgases. In 1804 he made balloonascensions to studymagnetic forces and toobserve the compositionand temperature of the airat different altitudes.
  25. 25. #3. Gay-Lussac’s Law - 1802•The pressure and Kelvin temperature ofa gas are directly proportional, providedthat the volume remains constant. P P2 1 = T1 T2•How does a pressure cooker affect the timeneeded to cook food? (Note page 422)•Sample Problem 14.3, page 423
  26. 26. #4. The Combined Gas LawThe combined gas law expresses therelationship between pressure, volumeand temperature of a fixed amount ofgas. PV1 P2V2 1 = T1 T2Sample Problem 14.4, page 424
  27. 27.  The combined gas law contains all the other gas laws! If the temperature remains constant... P1 x V1 P2 x V2 = T1 T2 Boyle’s Law
  28. 28.  The combined gas law contains all the other gas laws! If the pressure remains constant... P1 x V1 P2 x V2 = T1 T2 Charles’s Law
  29. 29. The combined gas law containsall the other gas laws!If the volume remains constant...P1 x V1 P2 x V2 = T1 T2 Gay-Lussac’s Law
  30. 30. Section 14.3 Ideal Gases OBJECTIVES: Compute the value of an unknown using the ideal gas law.
  31. 31. Section 14.3 Ideal Gases OBJECTIVES: Compare and contrast real an ideal gases.
  32. 32. 5. The Ideal Gas Law #1 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) The other units must match the value of the constant, in order to cancel out. The value of R could change, if other units of measurement are used for the other values (namely pressure
  33. 33. The Ideal Gas Law We now have a new way to count moles (the amount of matter), by measuring T, P, and V. We aren’t restricted to only STP conditions: PxV n= RxT
  34. 34. 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!
  35. 35. Ideal Gases There are no gases for which this is true (acting “ideal”); however, Real gases behave this way at a) high temperature, and b) low pressure. Because at these conditions, a gas will stay a gas Sample Problem
  36. 36. #6. Ideal Gas Law 2Equation: P x V = mxRxT M Allows LOTS of calculations, and some new items are: m = mass, in grams M = molar mass, in g/mol Molar mass = m R T PV
  37. 37. Density Density is mass divided by volume m D= Vso, m MP D= = V RT
  38. 38. Ideal Gases don’t exist, because:1. Molecules do take up space2. There are attractive forces between particles - otherwise there would be no liquids formed
  39. 39. 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
  40. 40. 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.
  41. 41. Section 14.4Gases: Mixtures and Movements OBJECTIVES: Relate the total pressure of a mixture of gases to the partial pressures of the component gases.
  42. 42. Section 14.4Gases: Mixtures and Movements OBJECTIVES: Explain how the molar mass of a gas affects the rate at which the gas diffuses and effuses.
  43. 43. #7 Dalton’s Law of Partial Pressures For a mixture of gases in a container, PTotal = P1 + P2 + P3 + . . .•P1 represents the “partial pressure”,or the contribution by that gas.•Dalton’s Law is particularly useful incalculating the pressure of gasescollected over water.
  44. 44. Connectedto gasgenerator Collecting a gas over water
  45. 45.  If the first three containers are all put into the fourth, we can find the pressure in that container by adding up the pressure in the first 3: 2 atm + 1 atm + 3 atm = 6 atm 1 2 3 4
  46. 46. 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.
  47. 47. •Diffusion:describes the mixingof gases. The rateof diffusion is therate of gas mixing.•Molecules movefrom areas of highconcentration to lowconcentration.
  48. 48. Effusion: a gas escapes through a tinyhole in its container -Think of a nail in your car tire… Diffusion and effusion are explained by the next gas law: Graham’s
  49. 49. 8. Graham’s Law RateA √ MassB = RateB √ MassA The rate of effusion and diffusion is inversely proportional to the square root of the molar mass of the molecules. Derived from: Kinetic energy = 1/2 mv2 m = the molar mass, and v = the velocity.
  50. 50. Graham’s Law With effusion and diffusion, the type of particle is important:  Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass. Helium effuses and diffuses faster than nitrogen – thus, helium escapes from a balloon quicker than many other gases