Chapter 12: States Of Matter

1. Chapter 12: States Of Matter Sec. 12.1: Gases

2. Objectives • Use the kinetic-molecular theory to explain the behavior of gases. • Describe how mass effects rates of diffusion and effusion. • Explain how gas pressure is measured and calculate the partial pressure of a gas.

3. Properties of Substances • Chemical & physical properties of substances depend on composition (the atoms present) & structure (arrangement of atoms). • However, substances that are gases display similar properties despite different compositions!

4. Kinetic Molecular Theory (1860) • Gases studied were molecules. • Objects in motion have energy called kinetic energy. • The kinetic molecular theory describes the behavior of gases in terms of particles in motion.

5. Kinetic Molecular Theory • The kinetic molecular theory assumes that gas particles have a VERY SMALL volume and that they are separated from one another by a LARGE volume of space. • Because they are so far apart, there is no attraction or repulsion between gas particles.

6. Kinetic Molecular Theory • Gas particles are in constant, random motion. • They move in straight lines until collision. • Collisions between gas particles are elastic. (There is no overall loss of kinetic energy.)

7. Kinetic Molecular Theory • 2 factors determine the kinetic energy of a gas particle: mass and velocity Within a gas sample, the mass does not vary but velocity will. Therefore, when we talk about KE, we really mean average KE.

8. Kinetic Molecular Theory • Temperature is a measure of the average kinetic energy of the particles in a sample of matter. • At a given temperature, all gas particles will have the SAME average kinetic energy.

9. Behavior of Gases • The constant motion of gas particles allows a gas to expand until it fills its container.

10. Behavior of Gases • Gases have a low density. (Remember: D = mass/volume) • There are fewer gas particles in a given volume than in the same volume of a liquid or solid. • A great deal of space exists between the gas particles.

11. Behavior of Gases • Gases are compressible (able to have their volume reduced) because there is so much empty space between gas particles.

12. Behavior of Gases • Diffusion is the term used to describe the movement of one material through another. Gases have no forces of attraction for one another so diffusion is possible. • Due to diffusion, gas particles tend to move from areas of high concentration to areas of low concentration, until they are evenly distributed.

13. Behavior of Gases • Rate of diffusion depends on the mass of the gas particles. • Light particles, at the same temperature as heavier particles, will have a greater velocity. They will therefore diffuse quicker. • Effusion is related to diffusion. During effusion, a gas escapes through a tiny opening.

14. Behavior of Gases • Graham’s law of effusion states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass. This law can also be applied to diffusion rates.

15. Practice Problems • What is the ratio of the diffusion rate of ammonia to hydrogen chloride? • Calculate the ratio of effusion for neon to nitrogen. • Calculate the ratio of diffusion rates for carbon monoxide to carbon dioxide.

16. Gas Pressure • Pressure is defined as force per unit area. • Gas particles exert pressure when they collide with the walls of their container.

17. Gas Pressure • Since pressure is a result of collisions between all of the gas particles and the surfaces around them, the amount of pressure increases when the number of particles in a given volume increases.

18. Atmospheric Pressure • The gas particles in air move in all directions, and so, exert air pressure in all directions. • There is less air pressure at high altitudes because there are fewer particles present, since the force of gravity is less. • Torricelli was the first to demonstrate that air exerted pressure. • He invented the barometer, an instrument used to measure atmospheric air pressure.

19. Atmospheric Pressure • A barometer has a closed tube that is inverted in a pool of Hg. The Hg rises & falls in the tube in response to the amount of air pressure applied to the Hg.

20. Atmospheric Pressure • Torricelli showed that at the Earth’s surface, the height of the Hg in the barometer was always about 760mm Hg. (mm Hg stands for millimeters of mercury). • This is considered standard air pressure.

21. Units of Pressure • The SI unit of pressure is the pascal (Pa). Standard air pressure is 101,300 Pa or 101.3 kPa. • Standard air pressure in more traditional units is: • 14.7 psi (pounds per square inch) • 760 torr (1 torr = 1 mm Hg) • 1 atm (atmosphere)

22. Practice Problems • Determine the value of each in kPa. • 3.5 atm • 930 torr • 560 mm Hg

23. Measuring Gas Pressure • A closed or open ended manometer is used to measure gas pressure in a closed container. • In a manometer, the difference in the levels levels of Hg in the U-tube is used to calculate the gas pressure in mm Hg.

24. Dalton’s Law of Partial Pressure • This law states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. • The portion of the total pressure contributed by a single gas is called the partial pressure.

25. Dalton’s Law of Partial Pressure • Mathematically: P1 + P2 + P3 + …. = PT We add the pressure of each gas in a mixture. Their sum is equal to the total pressure of gas in the container.

26. Practice Problems • A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of oxygen if the partial pressure of carbon dioxide is 0.70 atm and that of nitrogen is 0.12 atm? • Find the total pressure for a mixture that contains 4 gases with partial pressures of 5.00 kPa, 4.56 kPa, 3.02 kPa, and 1.20 kPa.