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.)
An elastic collision is one in which no kinetic energy is lost.

7. Kinetic Molecular Theory
2 factors determine the kinetic energy of a gas particle: mass and velocity
KE = kinetic energy
M = the mass of the particle
V = the particle’s velocity
In a sample of a single gas, all particles have the same mass but all particles do not have the same amount of kinetic energy.
Kinetic energy and temperature are related.
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.
Unlike solids and liquids, there is much empty space between the particles in a gas. Applying a force pushes the particles into the empty space, compressing the gas.

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.
It also applies to rates of diffusion.

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.
1.47
Ammonia molecules diffuse about 1.5 times as fast as hydrogen chloride molecules. This ratio is logical because molecules of ammonia are about half as massive as molecules of hydrogen chloride.

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.
(a) The mercury levels are equal inside and outside the open-end tube because the tube is open to the atmosphere and filled with air. (b) A column of mercury 760 mm high is maintained in the closed-end tube for standard atmospheric pressure. The space above the mercury is devoid of air (a vacuum), containing only a tiny trace of mercury vapor

20. Atmospheric Pressure
Torricelli showed that at the Earth’s surface, the height of the Hg in the barometer was always about 760 mm 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 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.