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Gases

This chapter explores the properties of gases, including potential and kinetic energy, as well as the Kinetic-Molecular Theory and Gas Laws such as Boyle's Law, Gay-Lussac's Law, Charles' Law, and Avogadro's Law.

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Gases

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  1. Gases Chapter 13

  2. Solids, liquids, and Gases Compare the position and motion of the three states of matter. Chapter 13

  3. Energy • Potential Energy • Stored energy – due to position • Particles are attracted to one another. More energy is required to keep particles farther apart. • Which of the 3 states has the highest potential energy? • Kinetic Energy • Motion energy – related to temperature • The faster the particles are moving, the higher the kinetic energy, the higher the temperature (average kinetic energy) • Which of the 3 states has the highest kinetic energy? Chapter 13

  4. Kinetic-Molecular Theory • Theory developed to explain gas behavior • To describe the behavior of a gas, we must first describe what a gas is: • Gases consist of very small particles each of which have a mass. • The distance between gas particles are relatively large. Volume of individual molecules is negligible compared to volume of container. • Gas particles in rapid, constant, random motion. Chapter 13

  5. Kinetic-Molecular Theory (Cont’d) • Collisions between gas particles are perfectly elastic. • Energy can be transferred between molecules, but total kinetic energy is constant at constant temperature. • No energy is lost during collisions. Chapter 13

  6. Kinetic-Molecular Theory (Cont’d) The average kinetic energy of gas particles depends only on the temperature of the gas. What happens as temperature increases? Gas particles exert no forces on each other. Intermolecular forces (forces between gas molecules) are negligible. • Which statements in KMT are assumptions? Chapter 13

  7. Let’s generate some gas Properties of a gas • Expand to fill a volume (expandability) • Compressible • Takes shape of container • Diffuses and flows Chapter 13

  8. Variables that can be measured for gases • Temperature • Volume • Amount • Pressure Chapter 13

  9. Temperature (T) • Measured in Fahrenheit, Celsius, or Kelvin. • For this chapter, we have to use Kelvin. ◦C = K – 273 K = ◦C + 273 Volume (V) • Measured in Liters, cubic meters, gallons, etc… Amount (n) • Measured in moles Chapter 13

  10. Pressure (P) – what causes pressure? 1 atm = 760 mmHg = 760 torr = 101.3 kPa The pressure of a gas is measured using a manometer. *Atomspheric pressure is measured using a Barometer. Chapter 13

  11. Unit Conversions 22◦C = ______ K 37 K = ______ ◦C 18 mL = ______ L 4.3 L = ______ cm3 500 cm3= ______ mL 2.2 dm3 = _______ L Chapter 13

  12. Unit Conversions 2.8 g N2 = ______ mol 612 mol SO3 = _______ g 22 kPa = ______ atm 289 mmHg = _______ kPa 4.3 atm = ______ torr 518 kPa = _______ mmHg Chapter 13 12

  13. The Gas Laws • There are four variables required to describe a gas: • Amount of substance: moles (n) • Volume of substance: volume (V) • Pressures of substance: pressure (P) • Temperature of substance: temperature (T) • The gas laws will hold two of the variables constant and see how the other two vary (n,V,P,T) Chapter 13

  14. Today’s Lab – Boyle’s Law • We will maintain a constant temperature and number of moles of gas. • So we will vary the Pressure and the Volume and see how they relate. PV = k or P/V = k inverse direct Chapter 13

  15. Bolyle’s Law Lab – 16 pts total • Heading – 1 pt • Purpose – 1 pt • Procedure – 1 pt • Data – 4 pts (make sure they have units for the 3rd and 4th columns. • Graph – 3 pts (should have a title and labeled axis) • Questions – 6 pts (in complete sentences) 1) 1/2 2) doubled 3) inverse 4) PV=k 5) The pressure and volume of a gas at a constant temperature are inversely proportional to each other. 6) Source of error. Chapter 13

  16. Variables for Gases Discussed Before • T – Temperature • V – Volume • P – Pressure • n – Amount of a substance (moles) Chapter 13

  17. Boyle’s Law – peeps in bell jar demo The Pressure-Volume Relationship Boyle’s Law - The volume of a fixed quantity of gas is inversely proportional to its pressure at a constant temperature. Chapter 13

  18. A gas occupies 22 L at 2.43 atm. What is the new volume if the pressure changed to 5.11 atm? Chapter 13

  19. Gay-Lussac’s Law – can crush demo The Pressure-Temperature Relationship: Gay-Lussac’s Law – As the temperature of an enclosed gas increases, the pressure increases if the volume is constant. Chapter 13

  20. A gas has a pressure of 1.47 atm at 303 K. What is the new temperature if the pressure changed to 680 mmHg? Chapter 13

  21. Charles Law – ivory soap demo The Temperature-Volume Relationship: Charles’s Law - the volume of a fixed quantity of gas at constant pressure increases as the temperature increases. Chapter 13

  22. A gas occupies 14 L at 275 K. What is the new volume if the temperature changed to 297 K? Chapter 13

  23. Avogadro’s Law The Quantity-Volume Relationship: Avogadro’s Law - The volume of gas at a given temperature and pressure is directly proportional to the number of moles of gas. Chapter 13

  24. A balloon contains 1.98 mol of a gas and has a volume of 4.2 L. Some of the gas was let out to give a volume of 3.1 L. What is the amount of gas left in the balloon? Chapter 13

  25. The Gas Laws Summary Boyle’s Law Charles’ Law Avogadro’s Law Guy-Lussac’s Law Chapter 13

  26. The Ideal Gas Equation Combine the gas laws (Boyle, Charles, Guy-Lussac, Avogadro) yields a new law or equation. Ideal gas equation: PV = nRT P = pressure (atm or mmHg or kPa) V = volume (L) n = amount (mol) R = gas constant T = temperature (K) Chapter 13

  27. Finding R with Dry Ice Lab PV = nRT What units are used for R? Chapter 13

  28. Calculating R with Dry Ice Lab – 20 pts total • Heading – 1 pt • Purpose – 1 pt • Procedure – 1 pt • Data – 10 pts (make sure they have units) • Calculations – 2pts • Conclusion – 5 pts • Their value for R • Literature value for R • % error • Sources of error • How they affected their results (higher or lower and must be consistent with their result for R) Chapter 13

  29. Calculating R We define STP (standard temperature and pressure) as 0C = 273 K, and 1 atm = 760 mmHg = 101.3 kPa Volume of 1 mol of gas at STP is 22.4 L. PV = nRT Chapter 13

  30. The Ideal Gas Equation-Finding R Chapter 13

  31. Combined Gas Law • So in cases with changing conditions and Chapter 13

  32. Ideal vs. Combined Gas Law-When Use • Ideal Gas Law: Conditions not changing • Combined Gas Law: Conditions Changing Chapter 13

  33. Example 1 • How many moles of a gas at 100 degrees C does it take to fill a 1.0 L flask to a pressure of 1.50 atm? Not a changing situation so use Ideal Gas Law PV=nRT Chapter 13

  34. Example 2 At 60 Celsius a 0.10 L sample of a gas has a pressure of 75.6 kPa. What would its volume be at STP? Changing situation so use combined gas law. Chapter 13

  35. Example 3 • What pressure would 3.55 grams of argon gas be under in a 2.40 L cylinder at -35 Celsius? Not a changing situation so use Ideal Gas Law PV=nRT Chapter 13

  36. Example 4 • What is the density of bromine gas the gas fills a 52.5 L cylinder at 145 K and 583 mmHg? Not a changing situation so use Ideal Gas Law PV=nRT Chapter 13

  37. Example 5 A gas occupies 4430 mL at 30 Celsius. It was transferred in a 3.5 L cylinder. What is the new temperature. Changing situation so use combined gas law. Chapter 13

  38. Demos Egg in E-flask Fountain Chapter 13 38

  39. Tank Car • This tank car was cleaned with steam then all the valves were shut and tank car was sealed. The workers went home and when they came back the next morning this is what they saw. Chapter 13

  40. Chapter 13

  41. Chapter 13

  42. Chapter 13

  43. Ideal gases behave “ideally” according to the kinetic molecular theory and follow the ideal gas law PV = nRT. Ideal vs. Real gases But kinetic molecular theory has several assumptions that work most of the time but not always. Chapter 13

  44. 1) The volume of gas particles are so small and the spaces between particles are so large. Therefore KMT assumes that gas particles have no volume. Kinetic Molecular Theory Assumptions 2) The distance between gas particles are very large. Therefore KMT assumes that there is no attraction (IMF) between gas particles. Chapter 13

  45. Under most conditions the volume of gas particles is negligible. But Real gases have volume But at small volumes or if the gas particles are large, the volume of the gas particles become significant. To compensate, we plug in the following equation for V in PV = nRT Chapter 13

  46. Under most conditions, the intermolecular forces between gas particles are negligible. And real gases have attractive forces (IMF) But at high pressure and low temperatures, the attractive forces become significant. To compensate for the pressure difference caused by IMF we plug in the following equation instead of P in PV = nRT Chapter 13

  47. Ideal vs. Real gases Which gas behaves more ideally? Ne or HCl? Neon because it’s particles have a smaller volume and weaker intermolecular forces. Chapter 13 47

  48. In PV = nRT , plugging in Ideal vs. Real gases We get the Van der Waals Equation Chapter 13

  49. Example • Use the Van der Waals equation to calculate the temperature of 3.6 moles of nitrogen gas in a 4.6 L cylinder at 2.5 atm if a and b for nitrogen are 1.390 and 0.03910 respectively. Chapter 13

  50. Gas Mixtures and Partial Pressures Dalton’s Law Dalton’s Law - In a gas mixture the total pressure is given by the sum of partial pressures of each component: Pt = P1 + P2 + P3 + … - The pressure due to an individual gas is called a partial pressure. Chapter 13

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