Presentation on theme: "Kinetic Molecular Theory of Gases and the Gas Laws Mr. Nelson Chemistry."— Presentation transcript:
Kinetic Molecular Theory of Gases and the Gas Laws Mr. Nelson Chemistry
Properties of Gases Gases are fluids Fluids are any substance that flows Gases are highly compressible Example: Tire pressure Gases completely fill containers Gases have lower densities than liquids and solids
Kinetic Molecular Theory KMT describes the motion of the particles Particles have the same motion as billiard balls http://intro.chem.okstate.edu/NSFCCLI/GasLa w/GLP.htm
Kinetic Molecular Theory of Gases Assumptions: Gas molecules are in constant, random motion Gas molecules are separated by large distances Gas molecules have no attractive/repulsive forces
Temperature of Gases Temperature and energy of gases are directly proportional As the temperature increases, kinetic energy of the molecules increases As temperature decreases, kinetic energy will also decrease
Pressure of Gases At sea level, the standard gas pressure is 1 atmosphere Pressure is the force exerted by gas molecules Standard Temperature and Pressure (STP) is equal to 1 atm and 0 °C
Converting Pressure Example Convert 72.7 atmospheres (atm) into kilopascals (kPa)
The Gas Laws Variables in Gas Equations: P = Pressure (kPa or atm) V = Volume (L) T = Temperature (K) n = amount of gas (moles)
Boyles Law States that for a fixed amount of gas at constant temperature the volume of the gas is inversely proportional to the pressure of a gas Pressure Volume
Boyles Law Example Problem The pressure on 2.50 L of anesthetic gas changes from 105 kPa to 40.5 kPa. What will be the new volume if the temperature remains constant?
Boyles Law Example Problem A high-altitude balloon contains 30.0 L of helium gas at 103 kPa. As the balloon rises, you record a new volume of 35.0 L. What is the atmospheric pressure in kPa? (Assume constant temperature)
Charless Law States that the volume of a gas is directly proportional to the Kelvin temperature if the pressure remains constant Volume Temperature
Charless Law Example Problem The air in a hot air balloon has a volume of 400.0 L at 30.0°C (303 K). What will the volume be if the temperature is raised to 120.0 °C (393 K)?
Charless Law Example Problem An aerosol can has a volume of 3.00 x 10 2 mL at 150.0°C is heated until its volume is 6.00 x 10 2 mL. What is the new temperature (in K) of the gas if pressure remains constant?
Gay-Lussacs Law States that the pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant Pressure Temperature
Gay-Lussacs Law Example Problem The gas left in a used aerosol can is at a pressure of 103 kPa at 25 ° C. If this can is thrown onto a fire, what is the pressure of the gas when its temperature reaches 928 ° C?
Gay-Lussacs Law Example Problem A sealed cylinder of gas contains nitrogen gas at 1.00 x 10 3 kPa pressure and a temperature of 20.0 ° C. The cylinder is left in the sun, and the temperature of the gas increases to 50.0 ° C. What is the new pressure in the cylinder?
Combined Gas Law A single equation that combines all the gas laws:
Combined Gas Law Example Problem A gas takes up a volume of 17 liters, has a pressure of 2.3 atm, and a temperature of 299 K. If I raise the temperature to 350 K and lower the pressure to 1.5 atm, what is the new volume of the gas?
Ideal Gas Law Relates the gas laws and the amount of gas Requires the gas constant, R R can be a different number depending on the units
PV = nRT Example Problem A container of 3.0 L of nitrogen (N 2 ) is at a pressure of 4.5 x 10 2 kPa and a temperature of 39 ° C. How many grams of N 2 are in the container?
Ideal Gas Law Example Problem What pressure will be exerted by 0.450 mol of a gas at 25.0 °C if it is contained in a 0.650 L vessel?
Avogadros Hypothesis Equal volumes of gases at the same temperature and pressure contain equal numbers of particles Due mainly to the large amount of empty space between particles From this, scientists have determined that 1 mol = 22.4 L at STP
This was not well accepted Why? Tennis balls vs. Bowling balls