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Gas Laws

Gas Laws. Read pp. 330-340. 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).

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Gas Laws

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  1. Gas Laws Read pp. 330-340

  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 • 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.

  5. Gas Pressure • Pressure is the force per unit area. • Air is a mixture of N2 gas, O2 gas and small amounts of CO2 gas. Each gas exerts its own pressure, which contributes to the total pressure of air = air pressure. • Air pressure changes when altitude changes

  6. Measuring Pressure • Pascal (Pa) and kiloPascal (kPa), where 1 kPa is 1000 Pa • Atmospheric pressure(atm) measure of the pressure exerted by air • Air is typically measured with 2 units: • kPa • atm (1 atm = 101.3 kPa) • Note: Standard Ambient Temperature and Pressure (SATP) is 25C & 100 kPa

  7. Boyle’s Law • The volume of a gas ↓ as the pressure ↑. • Equation: p1v1 = p2v2 Explains why you feel your chest tighten as you go deeper in the water or be at high altitudes (where the pressure is greater).

  8. Example 1 • Initial V = 43 L at 1.04 atm Final V = ? at 2.94 atm

  9. Example 2 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?

  10. Charles’ Law • As the temperature of a gas ↑, the volume also ↑. Assume pressure is constant. • Equation:

  11. Measuring Temperature • When doing calculations with Charles’ Law, you must convert the temperature from Celsius (C) to Kelvin (K) • 0°C = 273K • If temperature is in °C, just add 273 to it.

  12. Example 1 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?

  13. Example 2 A sample of neon has a volume of 1.83 L at 23.5°C. At what temperature would the gas occupy 5.00 L?  Assume pressure is constant.

  14. 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: Ptotal = P1 + P2 + P3 ….

  15. Example 1 A scuba diver’s air tank contains a mixture of nitrogen gas and oxygen gas. The partial pressure for N is 110 atm and for O, it is 28 atm. What is the total partial pressure?

  16. Example 2 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?

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