CHAPTER 14 THE BEHAVIOR OF GASES: 14.1 Properties of gases 14.2 The Gas Laws 14.3 Ideal Gases

1. CHAPTER 14 THE BEHAVIOR OF GASES: 14.1 Properties of gases 14.2 The Gas Laws 14.3 Ideal Gases 14.4 Gases: Mixtures and Movements

2. 14.1 Properties of Gases • Why are gases easier to compress than solids? • What are the three factors affect gas pressure?

3. Compressibility: • Gases expand to fill their containers • Gases can be compressed into smaller volumes • Compressibility – is a measure of how much the volume of matter decreases under pressure. • Gases are easily compressed because of the space between the particles. • At room temperature, the distance between particles in an enclosed gas is about 10 times the diameter of a particle.

4. Factors Affecting Gas Pressure: -Four variables are used to describe a gas. a. Pressure in kilopascals (P) b. Volume in Liters (V) c. Temperature in Kelvins (T) d. Number of moles (n)

5. Why a Curved Bottom? In most aerosol cans, the bottom curves inward. This serves two functions:

6. 14.3 Ideal Gases: • -What is needed to calculate the amount of gas in a sample at given conditions of volume, temperature, and pressure? • Under what conditions are real gases most likely to differ from ideal gases?

7. Ideal Gas Law: • The combined gas law can be modified to include the number of moles. • You can introduce moles into the combined gas law by dividing each side of the equation by n. • The equation (P x V)/(T x n) is a constant for ideal gases that conform to the gas laws. • The constant has the symbol (R) such that R = (P x V)/(T x n) • The value of R is 8.31(LkPa/Kmol) • The gas law includes all four variables • PV = nRT

8. Ideal Gas and Real Gas -An ideal gas is one that follows the gas laws at all conditions of pressure and temperature, precisely to the kinetic theory. -the particles can have no volume and no attraction to other particles in the gas. -There is no gas at which these assumptions are true. -However, real gases behave very much like an ideal gas at many conditions of temperature and pressure. -Real gases differ most from an ideal gas at low temperatures and high pressure. -Real gases deviate from the ideal a. When the ratio is greater than 1 and the curve rises above the ideal gas line. Molecular volume dominates due to more kinetic energy b. When the ratio is less than 1 and the curve drops below the line. Intermolecular attractions dominate.

9. Liquid Gas? Since the product is liquid at room temperature, it is simply poured in before the can is sealed. The propellant, on the other hand, must be pumped in under high pressure after the can is sealed. When the propellent is kept under high enough pressure, it doesn't have any room to expand into a gas. It stays in liquid form as long as the pressure is maintained. (This is the same principle used in a liquid propane grill.)

10. When the valve is open, the pressure on the liquid propellant is instantly reduced. With less pressure, it can begin to boil. Particles break free, forming a gas layer at the top of the can. This pressurized gas layer pushes the liquid product, as well as some of the liquid propellant, up the tube to the nozzle.

11. When the liquid flows through the nozzle, the propellant rapidly expands into gas. In some aerosol cans, this action helps to atomize the product, forming an extremely fine spray. In other designs, the evaporating propellant forms bubbles in the product, creating a foam. The consistency of the expelled product depends on several factors, including: The chemical makeup of the propellant and product The ratio of propellant to product The pressure of the propellant The size and shape of the valve system

12. Aerosol can and high temperature

13. 14.2 The Gas Laws -How are the pressure, volume, and temperature of a gas related? -When is the combined gas law used to solve problems? Important Concepts: -Boyle’s Law -Charles’s Law -Gay-Lussac’s Law -Combined Gas Law

14. Boyle’s Law: Pressure and Volume For a given mass of a gas at constant temperature, the volume of the gas is inversely proportional to the pressure. P1 x V1 = P2 x V2 The graph of an inverse relationship is always a curve.

15. Charles’s Law: Temperature and Volume • The volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure remains constant. V1 / T1 = V2 / T2 The ratio of the variables is always constant in a direct relationship and the graph is always a straight line

16. Gay-Lussac’s Law: Pressure and Temperature -The pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant. P1 / T1 = P2 / T2

17. The Combined Gas Law - Describes the relationship among the pressure, temperature, and volume of an enclosed gas. P1 x V1 / T1 = P2 x V2 / T2 By remembering the combined gas law, all other laws can be derived as long as one variable remains constant.

18. 14.3 Ideal Gases • What is needed to calculate the amount of gas in a sample at given conditions of volume, temperature, and pressure? • Under what conditions are teal gases most likely to differ from ideal gases?

19. Ideal Gas Law: • -The combined gas law assumes that the amount of gas does not vary. • Useful to calculate the number of moles of a contained gas. • The number of moles of a gas is directly proportional to the number of particles. • The volume occupied by a gas at a specified temperature and pressure • also must depend on the number of particles. • -Moles are directly proportional to volume too. • - Moles is introduced to the equation as follows; • P1 x V1 / T1 x n1 = P2 x V2 / T2 x n2 • The equation shows that (PxV)/(Txn) is a constant. • This holds for ideal gases – gases that conform to the gas laws. • The value of the constant can be calculated to be 8.31(LkPa)/(Kmol) • This is known as the ideal gas constant (R) • The ideal gas law has four variables • PV = nRT

20. Ideal Gases and Real Gases: • An ideal gas must conform to the kinetic theory such that its particles could have no volume, and no attraction between particles in a gas • So a true “ideal gas” does not exist. • Rather, at many conditions of temperature and pressure, real gases behave like an ideal gas. • Because real gases do have volume and an attraction between particles is present – a gas can condense or solidify when compressed or cooled. • Real gases differ most from an ideal gas at low temperatures and high pressures.

21. -For the ratio PV/nRT as pressure increases • For an ideal gas, the result is a horizontal line because the ratio is • always 1 • -For real gases at high pressure, the ratio may deviate from the ideal • When the ratio is greater than 1, the curve is above the line and • molecular volume dominates • When the ratio is less than • 1, the curve is below the line • and intermolecular • attractions dominate.

22. 14.4 Gases: Mixtures and Movements -How is the total pressure of a mixture of gases related to the partial pressures of the component gases? -How does the molar mass of a gas affect the rate at which the gas effuses or diffuses?

23. Dalton’s Law: • Gas pressure results from collision of particles in a gas with its • container. • If the number of particles in a given volume increase, pressure • increases. • b. If the average kinetic energy of particles increases in a given • volume, pressure increases. • -Particles in a mixture of gases at the same temperature have the • same average kinetic energy. Particle type is not important. • -Partial Pressure: is the pressure contribution of each gas in • the mixture. • -Total Pressure: is the sum of the partial pressures in the gas mixture • -Dalton’s law of partial pressures: At constant volume and temperature • the total pressure exerted by a mixture of gases is equal to the sum • of the partial pressures of the component gases

24. Graham’s Law: • -Diffusion: is the tendency of molecules to move toward areas of • lower concentration until the concentration is uniform. • -Effusion: A gas escapes through a tiny hole in its container. • -Gases of lower molar mass diffuse and effuse faster than gases • of higher molar mass • -Graham’s law of effusion: The rate of effusion of a gas is inversely • proportional to the square root of the gas’s molar mass. Also • applicable to diffusion. • Comparing rates of effusion • rate of effusion to the particles • speed and mass.