Gases
- Solids: they have a define shape and volume
- Liquids: define volume but assumes the shape of the container
- Gases: no definite shape or volume
- Gases flow like liquids and are fluids.
- The difference between a gas and a liquid is the distance between the molecules.
- barometer: a common instrument to measure air pressure
- Pressure Conversion:
- 1 atm = 760 mm Hg
- 1 atm = 760 torr
- 1 atm = 14.7 psi
- 1 atm = 101,325 Pa
- Atmospheric pressure results from the mass of the air being pulled toward the center of the earth by gravity.
- This is also known as weight of air.
Boyle's Law
- At constant temperature the volume occupied by a fixed amount of gas is inversely proportional to the pressure on the gas.
- Pressure and volume are inversely proportional. If one doubles, the other decreases by one-half.
- PV = PV
- Ex. When the number of the air molecule in the tire is increased, the pressure is increased.
- Assumption: Temperature must be consistent.
Charles's Law
- The French physicist Jacques Charles (1746-1823) showed that the volume of a given amount of gas at constant pressure will increase with the temperature of the gas.
- Temperature is proportional to volume.
- As temperature increase volume increase.
- As temperature decrease volume decrease.
- Vi / Ti = Vf / Tf
Combined Gas Law
- Assumption: Number of moles constant.
- This law combines both Boyle’s law and Charles’s law.
- Pi Vi / Ti = Pf Vf / Tf
Avogrado's Law
- There is a relationship between the volume of a gas and the number of molecules presents in the gas sample.
- The number of moles is proportional to volume.
- As number of moles increase volume increase.
- As number of moles decrease volume decreases.
- Vi / ni = Vf / nf
Ideal Gas Law
- PV = nRT
- R = universal gas constant
- 0.08206 L*atm/(mol *K)
- Ideal gas law was derived from the observation of Boyle’s, Charles’s and Avogradro’s.
- Assumption: The gas must be ideal.
- Kinetic molecular theory of gases is a relatively simple model that attempts to explain the behavior of an ideal gas.
- Gases consist of tiny particles (atoms or molecules).
- These particles are so small, compared with the distances between them, that the volume(size) of the individual particles can be assumed to be zero.
- The particles are in constant random motion, colliding with the walls of the container. These collision with the walls cause the pressure exerted by the gas.
- The particles are assumed not to attract or to repel each other.
- The average kinetic energy of the gas particles is directly proportional to the Kelvin temperature of the gas.
Dalton's Law of Partial Pressure
- Many important gases contain a mixture of components.
- John Dalton was one of the first scientist to studied mixtures of gases.
- His observation became Dalton’s law of partial pressures.
- For a mixture of gases in a container, the total pressure exerted is the sum of the partial pressures of the gases present.
- P(total) = P1 + P2 + P3
- Pressure of a gas is not based on the element of the gas but the number of molecules of the gas.