1. Gas Laws

2. The States of Matter Recall: –Solids have a fixed, definite shape (strong forces between particles) –Liquids take the shape of its container (can be strong forces, but weaker than solids) –Gases have no shape or volume (weak or no forces between particles)

3. Kinetic Molecular Theory Explains the relationships between particles, the forces between them and the speed they move at. States: “All substances contain particles that are in constant, random motion.”

4. KMT and Gases Gases move at a far higher speed than solids/liquids and more randomly. They move continuously. Creates large spaces between particles  explains why gases are highly compressible Kinetic energy is the energy of movement. The faster the motion of an object, the greater the kinetic energy ↑ the temperature ↑ the particle speed = kinetic energy also ↑ as well. When heated, object takes up more space (expansion).

5. Gas Pressure Pressure is the force per unit area. Air is a mixture of N 2 gas, O 2 gas and small amounts of CO 2 gas. Each gas exerts its own pressure, which contributes to the total pressure of air = air pressure. –Air pressure changes when altitude changes As you go up towards the atmosphere, pressure ↓ As you go below water (beneath sea-level), pressure ↑

6. Measuring Pressure Pascal (Pa) and kiloPascal (kPa), where 1 kPa is 1000 Pa Atmospheric pressure (atm)  measure of the pressure exerted by air Units: 1.kPa 2.atm (1 atm = 101.3 kPa) 3.mmHg – millimeters of mercury (1 atm = 760 mmHg) Note: Standard Ambient Temperature and Pressure (SATP) is 25°C & 100 kPa

7. Conversions a) Convert 0.875 atm to mmHg b) Convert 98.35 kPa to atm a)1 atm = 0.875 atm 760 mmHg x x = 665 mmHg b) 1 atm = x. 101.3 kPa 98.35 kPa x = 0.970 atm

8. Air Pressure on Ground and 10,000 m in atmosphere On the ground, the air pressure is 1 atm (101 kPa). Up in the atmosphere, the pressure is less (<1 atm or <101 kPa) Gravity forces the air downward (air particles more dense on surface than up in atmosphere) Pressure ↓ with ↑ altitude

10. Boyle’s Law The volume of a gas ↓ as the pressure ↑. Equation: p 1 v 1 = p 2 v 2 Explains why you feel your chest tighten as you go deeper in the water (where the pressure is greater).

11. Atoms inside constantly colliding inside walls = exerting pressure As atmospheric pressure ↓, there are fewer air molecules to collide with on the outside So as pressure ↓ outside, the inside expands

12. Deep-Sea Diving When you deep-sea dive, the deeper you go underwater, the higher the pressure. In this case, pressure ↑ with ↓ altitude The increase in pressure causes your body’s organs to be compressed –Must take a slow rise to the top so your lungs don’t burst (it expands as you go back up). Go too fast, it’s just like opening a new bottle of soda (gases burst out) You are taking the weight of the water pushing down on you by gravity

13. Example The pressure exerted on a 240 mL sample of hydrogen gas at constant temperature is increased from 0.428 atm to 0.724 atm. What will the final volume of the sample be? P 1 V 1 = P 2 V 2 (0.428)(240) = (0.724)V 2 V 2 = 102.72 0.724 V 2 = 142 mL Take careful note of the units! On the test, you may have to convert units. Pressure units should be same on both sides of eq’n. Same for volume, etc.

14. Charles’ Law As the temperature of a gas ↑, the volume also ↑. Assume pressure is constant. Equation: V1 V2 T1 T2 =

15. Draw a flow chart to scientifically explain how a hot air balloon is lifted into the air and how it comes down to the ground.

16. Measuring Temperature When doing calculations with Charles’ Law, you must convert the temperature from Celsius (C) to Kelvin (K) 0°C = 273K Absolute ZERO is 0 K (equiv. to -273 °C). It is the lowest possible temperature. If temperature is in °C, just add 273 to it.

17. Example The volume of a gas inside a cylinder is 0.30 L at 25 ° C. The gas in the cylinder is heated to 315 ° C. What is the final volume of the gas after this heating? V 1 = V 2 T 1 T2 0.30 L = V 2 298 K 588 K V 2 = 0.59 L 25°C  298 K 315°C  588 K

18. Combined Gas Law Takes into account Boyle’s, Charles’ and Pressure-Temperature laws: Equation: P 1 V 1 = P 2 V 2 T 1 T 2

19. Example A toy balloon has an internal pressure of 1.05 atm and a volume of 5.0 L. If the temperature where the balloon is released is 200°C, what will happen to the volume when the balloon rises to an altitude where the pressure is 0.65 atm and the temperature is –150°C? 200°C  473 K –150°C  123 K P 1 V 1 = P 2 V 2 T 1 T 2 (1.05)(5) = (0.65)V 2 473123 V 2 = 2.1 L

20. Dalton’s Law of Partial Pressures The total pressure of a mixture of non-reacting gases is EQUAL to the sum of the partial pressures of the individual gases. Equation: P total = P 1 + P 2 + P 3 ….

21. Example A tank of compressed air that is used by a firefighter holds nitrogen at a partial pressure of 300 kPa. The tank has a total pressure of 385 kPa. What is the partial pressure of the oxygen in the tank? P total = P O2 + P N2 385 = P O2 + 300 P O2 = 85 kPa

22. Ideal Gas Law An “ideal” gas exhibits certain theoretical properties. Specifically, an ideal gas … –Obeys all of the gas laws under all conditions. –Does not condense into a liquid when cooled. –Shows perfectly straight lines when its V and T & P and T relationships are plotted on a graph. In reality, there are no gases that fit this definition perfectly. We assume that gases are ideal to simplify our calculations.

23. The Ideal Gas Law PV = nRT P = Pressure (in kPa)V = Volume (in L) T = Temperature (in K) n = moles R = 8.31 kPa L K mol R is constant. If we are given three of P, V, n, or T, we can solve for the unknown value. Recall, From Boyle’s Law: P 1 V 1 = P 2 V 2 or PV = constant From combined gas law: P 1 V 1 /T 1 = P 2 V 2 /T 2 or PV/T = constant R = 0.0821 atm L K mol

24. Example How many moles of H 2 is in a 3.1 L sample of H 2 measured at 300 kPa and 20°C? Givens: V = 3.1 L P = 300 kPa T = 20°C  293 K R = 8.31 kPa·L·K -1 ·mol -1 PV = nRT (300 kPA)(3.1 L) = n(8.31 kPa·L·K -1 ·mol -1 )(293 K) n = 0.382 mol