States of Matter: Gases

Unit Objectives Describe the properties of gases. Describe how pressure is measured Convert between various units of pressure Explain the relationships among pressure, volume, temperature, and amount of gas Use Boyle’s Law, Charles’s Law, Avogadro’s Law, Combined Gas, and Ideal Gas Laws to calculate changes in pressure, volume, temp. and amount of gas.

Unit Objectives Solve Problems using Dalton’s Law of partial pressures Describe the assumptions on which the kinetic molecular theory is based Calculate root mean square speeds of molecular gases List two molecular properties responsible for deviation from ideal gas behavior. Identify the conditions under which gases will be most likely to behave in a non-ideal manner Explain, compare and contrast diffusion and effusion of gases Calculate the pressure of a real gas as predicted by the van der Waals equation. Perform calculations for gases collected over water

Gaseous State of Matter has no distinct shape or volume so fills any container is easily compressed mixes completely with any other gas exerts pressure on its surroundings

Measuring Pressure barometer: measures atmospheric pressure invented by Torricelli, Italian scientist in 1643 glass tube filled with mercury is inverted in a dish mercury flows out of the tube until pressure of the Hg inside the tube is equal to the atmospheric pressure on the Hg in the dish

Measuring Pressure atmospheric pressure: results from mass of air being pulled toward the earth by gravity varies with altitude and weather conditions

Measuring Pressure manometer: measures pressure of gas in a container gas has less pressure than atmosphere if the Hg is closer to chamber gas has more pressure than atmosphere if the Hg is further from chamber gas pressure = atmospheric pressure ± h

Units of Pressure mmHg: most common since use Hg in manometers and barometers torr: equal to mmHg standard atmosphere (atm) Pascal (Pa): SI unit; equal to N/m2 1atm = 760mmHg = 760torr = 101,325Pa = 101.325kPa Example

States of Matter: Gases Part 2 Gas Laws

Boyle’s Law Discovered by Irish chemist, Robert Boyle Used a J-shaped tube to experiment with varying pressures in multistory home and effects on volume of enclosed gas P and V are inversely proportional PV = k holds precisely at very low pressures Example

Charles’ Law discovered by French physicist, Jacques Charles in 1787 first person to fill balloon with hydrogen gas and make solo balloon flight V and T are directly proportional V = kT

Charles’ Law for any gas, at -273.2°C, the volume is zero since negative volumes cannot exist, there cannot be a temperature lower than absolute zero (-273.2°C or 0 K) never actually been reached (0.000001 K has been) Kelvin system has no negative values Example

Avogadro’s Law Discovered by Italian chemistry, Avogadro in 1811 proposed that equal volumes of gases at the same temperature and pressure contain the same number of particles V = kn V and n are directly proportional Example

Gay-Lussac’s Law discovered in 1802 by Joseph Gay-Lussac P = kT P and T are directly proportional Example

Example: Pressure Conversions The pressure of a gas is measured as 49 torr. Represent this pressure in atmospheres, Pascals, and mmHg.

Example: Boyle’s Law Consider a 1.53-L sample of gaseous SO2 at a pressure of 5.6 x 103 Pa. If the pressure is changed to 1.5 x 104 Pa at constant temperature, what will be the new volume of the gas?

Example: Charles’ Law & Temp. A sample of gas at 15°C and 1 atm has a volume of 2.58 L. What volume will this gas occupy at 38°C and 1 atm?

Example: Avogadro’s Law Suppose we have a 12.2-L sample of gas containing 0.50 mol O2 at a pressure of 1 atm and temperature of 25°C. If all of this O2 were converted to O3 (ozone) at the same temperature and pressure, what would be the volume of O3?

Example: Gay-Lussac’s Law The gas in an aerosol can is at a pressure of 3.00 atm at 25°C. Directions on the can warn the user not to keep the can in a place where temperature exceeds 52°C. What would the gas pressure be in the can at 52°C?