1. Intro to Gases

2. First, remember the 3 states of matter…

3. Gases consist of small particles that are separated from one another by empty space. Gas: Particle size The volume of the particles is small compared with the volume of the empty space. Because gas particles are far apart, there are no significant attractive or repulsive forces among them.

4. Particles are in constant, random motion. Move in a straight line until they collide with other particles. Elastic Collision: no kinetic energy is lost (energy due to motion) Gas: Particle motion

5. Two factors determine energy: –Mass –Velocity In a sample of a single gas, all particles have the same mass but all particles do not have the same velocity. Therefore, all particles do not have the same kinetic energy. KE = ½ mv 2 Gas: Particle energy

6. Properties of Gases You can predict the behavior of gases based on the following properties: Pressure Volume Amount (moles) Temperature Lets review each of these briefly…

7. Pressure Pressure is defined as the force the gas exerts on a given area of the container in which it is contained. If you’ve ever inflated a tire, you’ve probably made a pressure measurement in pounds (force) per square inch (area).

8. Pressure Units SI unit of pressure is the Pascal (Pa). It is named for Blaise Pascal. One Pascal is equal to a force of one Newton per square meter: 1 Pa= 1 N/m 2

9. Pressure: Compressible Gasses can be compressed Compression – can be pushed into a smaller volume Expansion – can be pulled into a greater volume

10. Volume Volume is the three-dimensional space inside the container holding the gas. The SI unit for volume is the cubic meter, m 3. A more common and convenient unit is the liter, l. Think of a 2-liter bottle of soda to get an idea of how big a liter is. (OK, how big two of them are…)

11. Volume: Density Density – mass per unit of volume (g/cm 3 ) Density of gasses is much lower than the density of solids or liquids.

12. Temperature Temperature is a measure of the average kinetic energy of the particles in a sample of matter. Measured in Kelvin, K. The Kelvin scale starts at Absolute 0, which is -273.15°C. Temperature at which entropy reaches its minimum value (lowest temp possible). To convert Celsius to Kelvin, add 273.15.

13. Amount (moles) Amount of substance is tricky. As we’ve already learned, the SI unit for amount of substance is the mole, mol. Since we can’t count molecules, we can convert measured mass (in kg) to the number of moles, n, using the molecular or formula weight of the gas. By definition, one mole of a substance contains approximately 6.022 x 10 23 particles of the substance. You can understand why we use mass and moles!

14. Relationships Directly proportional – both variables will either increase or decrease together Inversely proportional – one variable will increase and the other variable will decrease

15. PTVPTV

16. Temperature & Pressure Temperature and pressure are directly proportional As one goes up, the other goes up As one goes down, the other goes down

17. Temperature & Volume Temperature and volume are directly proportional As one goes up, the other goes up As one goes down, the other goes down

18. Pressure & Volume Pressure and volume are Inversely proportional As one goes up, the other goes down As one goes down, the other goes up

19. Kinetic Molecular Theory 1.Particles of a gas are much smaller than the distance between particles, therefore, most of the volume of a gas is empty space. 2.The particles are in constant motion. The collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas. 3.There are no attractive or repulsive forces between the molecules. 4.The average kinetic energy of a collection of gas particles is related to the temperature of the gas. 5.There is a theoretical temperature at which the motion of the atoms/molecules stops, absolute zero!

20. Gas Law Calculations

21. Boyle’s Law –Robert Boyle –Studied relationship between pressure and volume when temperature was constant. – Determined they were inversely proportional. –P 1 V 1 = P 2 V 2

22. Charles’ Law –Jacques Charles – Studied relationship between volume and temperature when pressure is held constant – Determined they were directly proportional –V 1 T 2 = V 2 T 1

23. Gay-Lussac’s Law –Joseph Gay-Lussac –Studied relationship between pressure and temperature when volume is held constant –Determined they were directly proportional –P 1 T 2 = P 2 T 1

24. Gas Laws 1.Ideal Law 2.Combined Law 3.Dalton’s Law 4.Avogadro’s Law

25. Ideal Gas Law P = pressure, V = volume T = Temperature N = # of moles, R = 0.0821 PV = nRT

26. Combined Gas Law P = pressure, V = volume T = Temperature P 1 V 1 = P 2 V 2 or P 1 V 1 T 2 = P 2 V 2 T 1 T 1 T 2

27. Dalton’s Law The pressure of a mixture of gases is equal to the sum of the pressures of all of the constituent gases alone. P = pressure Pressure total = Pressure 1 + Pressure 2... Pressure n

28. Avogadro’s Law At a given temperature and pressure, equal volumes of gas contain equal numbers of moles. N = # of Moles, V = volume N 1 V 1 = N 2 V 2