GAS – state of matter that has NO DEFINITE VOLUME and NO DEFINITE SHAPE. Fig. 5: Arrangement of Particles in a Gas From : A Gas ALWAYS TAKES BOTH THE VOLUME AND THE SHAPE OF ANY CONTAINER INTO WHICH IT IS PLACED. If a gas is NOT in a container, it will spread out as far as it can.

Temperature Temperature is a measure of the average kinetic energy of matter. The greater the average kinetic energy, the greater the velocity of the particles of matter, the greater the temperature (and vice versa). The instrument used to measure the temperature (average kinetic energy) of matter is the thermometer.

1. Thermometer: A thin glass, capillary tube that contains a fluid (mercury (Hg)) that when heated will expand. When the fluid expands, it rises. This correlates to an increase in temperature. 2. Temperature Scales: There are several scales that are used to represent the temperature of substances. They include: Fahrenheit, Celsius, and Kelvin. Scientists most frequently use the Celsius and Kelvin scales.

Fahrenheit CelsiusKelvin 32 °0 ° °100 °373 Figure 9: Temperature Scales and the Boiling/Freezing Points of Water at Standard Conditions Boiling Point Freezing Point 0 Absolute Zero

Ex: #3 What Kelvin temperature is equal to 25°C? a) 248 K b) 298 K c) 100 K d) 200 K Conversion Formulas: C = CelsiusK = Kelvin 1. C = K – K = C + 273

VIII. Avogadro’s Hypothesis 1 mole 22.4 liters 6.02 x particles Gram Molecular (Formula) Mass Under the same conditions of temperature and pressure, equal volumes of all gases contain the same number of particles. For example, 1 liter of hydrogen gas will contain the same number of particles as 1 liter of oxygen gas.

IX. The Unique Properties of Gases a) Gases are greatly influenced by changes in temperature and pressure. b) The behavior of gases is due in part to the fact that the atoms/molecules that they are comprised of are greatly dispersed.

c) The Kinetic-Molecular Theory of Gases (Ideal Gases) 1. Gases are made from molecules that are in constant random motion. 2. Collisions between gas particles are completely elastic. 3. The volume of individual gas molecules is insignificant as compare to the overall volume of space that the gas occupies. 4. Individual gas molecules do not have attractive forces for each other. 5. The average kinetic energy of a gas is directly proportional to its temperature.

d) Exceptions to Ideal Gases 1. As opposed to an ideal gas, real gas molecules do have small but significant volumes as well as attractive forces. 2. These deviations become apparent under the conditions of low temperature and high pressure.

e) The Gas Laws 1.Boyle’s Law: an indirect relationship; states that at a constant temperature the volume of a gas decreases with increasing pressure. P 1 V 1 = P 2 V 2 Figure 16: Boyle’s Law From: le/SSGifs/26Fig2.gif KEY: P 1 = initial pressure P 2 = final pressure V 1 = initial volume V 2 = final volume

Example #1: Divers get “the bends” if they come up too fast because gas in their blood expands, forming bubbles in their blood. If a diver has 0.05 L of gas in his blood under a pressure of 250 atm, then rises instantaneously to a depth where his blood has a pressure of 50.0 atm, what will the volume of gas in his blood be? Do you think this will harm the diver? Example #2: Part of the reason that conventional explosives cause so much damage is that their detonation produces a strong shock wave that can knock things down. While using explosives to knock down a building, the shock wave can be so strong that 12 liters of gas will reach a pressure of 3.8 x 10 4 mm Hg. When the shock wave passes and the gas returns to a pressure of 760 mm Hg, what will the volume of that gas be?

2. Charles Law: a direct relationship; states that at a constant pressure, the volume of a gas will increase with a corresponding increase in the temperature (Kelvin). V 1 V 2 T 1 =T 2 Figure 17: Charles Law From: /GasLaw/charles.gif KEY V1 = initial volume T2 = final temperature V2 = final volume T1 = initial temperture

Example #1: The temperature inside my refrigerator is about 4 0 Celsius. If I place a balloon in my fridge that initially has a temperature of 22 0 C and a volume of 0.5 liters, what will be the volume of the balloon when it is fully cooled by my refrigerator? Example #2: A soda bottle is flexible enough that the volume of the bottle can change even without opening it. If you have an empty soda bottle (volume of 2 L) at room temperature (25 0 C), what will the new volume be if you put it in your freezer (-4 0 C)?

3.Combined Gas Law (P1) (V1) = (P2) (V2) T1 T2 NOTE: Standard Temperature and Pressure 0°C or 273 K kPa or 1 atm

Example #1: A gas that has a volume of 28 liters, a temperature of 45 0 C, and an unknown pressure has its volume increased to 34 liters and its temperature decreased to 35 0 C. If I measure the pressure after the change to be 2.0 atm, what was the original pressure of the gas?

4. Dalton’s Law of Partial Pressures: - Defines how the pressure, volume, and temperature of a gas depend upon the number of moles of the gas. P total = P 1 + P P n P t is the total pressure of a sample which contains a mixture of gases. P 1, P 2, P 3, etc. are the partial pressures (in the same units) of the gases in the mixture. The individual pressure of each gas (partial pressure) is directly related to it’s mole fraction. Example #1: A container holds three gases: oxygen, carbon dioxide, and helium. The partial pressures of the three gases are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container?

Example #2: A container holds three gases: oxygen, carbon dioxide, and helium. The partial pressures of the three gases are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container? Example #3: A tank contains grams of oxygen and grams of helium at a total pressure of 7.00 atmospheres. Calculate the following. a) How many moles of O 2 are in the tank? b) How many moles of He are in the tank? c) Total moles of gas in tank. d) Mole fraction of O 2. e) Mole fraction of He. f) Partial pressure of O 2. g) Partial pressure of He.

5. Gay-Lussac’s Law For a given amount of gas held at constant volume, the pressure is proportional to the absolute temperature. In a closed system, as temperature increases, the internal pressure exerted by a gas increases. P1=P2T1T2P1=P2T1T2 Example #1: A container is designed to hold a pressure of 2.5 atm. The volume of the container is 20.0 cm 3, and it is filled with air at room temperature (20°C) and normal atmospheric pressure. Would it be safe to throw the container into a fire where temperatures of 600°C would be reached?

6. Graham's Law of Diffusion/Effusion Diffusion - The rate at which two gases mix. Effusion - The rate at which a gas escapes through a pinhole into a vacuum. The rate at which gases diffuse is inversely proportional to the square root of their molar masses (MM). Rate 1 / Rate 2 = square root of (Mass 2 / Mass 1 )

7. Ideal Gas Law: P V = n R T P = pressure of the gas V = volume of the gas n = moles of gas T = temperature of the gas (Kelvin) R is the ideal, or universal, gas constant. R has the value J·K −1 ·mol −1 or L·atm·mol -1 ·K -1. The ideal gas law relates the variables of pressure, volume, temperature, and number of moles of gas within a closed system.