1-PROPERTIES OF GASES

Properties of Gases

  • Expansion: Gases expand to fill any available space.

  • Density: Gases have lower densities compared to liquids and solids of the same mass.

  • Color and Odor: Most gases are colorless and odorless.

  • Mixing: Gases can mix together in any proportion, forming a homogeneous solution as long as no reaction occurs.

  • Interrelation: The volume, temperature, and pressure of a gas are interrelated.

Gaseous State of Matter

  • Shape and Volume: Gases have no distinct shape or volume; they fill any container they occupy.

  • Compression: Gases are easily compressed.

  • Mixing: Gases mix completely with any other gas.

  • Pressure: Gases exert pressure on their surroundings.

Gases and the Kinetic Molecular Theory (KMT)

Kinetic Molecular Assumptions

  • Any gas sample consists mostly of empty space.

  • Particles move in rapid straight-line motion.

  • There is no loss of energy when particles collide.

  • In the gas phase, there are no attractive forces among particles.

  • At a given temperature, all gases have the same average kinetic energy.

    • Example: Samples of H2 and O2 at the same temperature possess the same kinetic energy.

Ideal and Real Gases

  • Ideal Gases: Follow all KMT assumptions.

  • Real Gases: Do not follow KMT completely; experience displacement.

    • At High Pressures: Forces of attraction become significant.

    • At Low Temperatures: Particle volume is a significant proportion of the total volume.

Measuring Pressure

Barometer

  • Measures atmospheric pressure.

  • Invented by Torricelli in 1643; uses a glass tube filled with mercury inverted in a dish.

  • Mercury flows until the pressure inside the tube equals atmospheric pressure.

Variables Affecting Barometer Measurements

  • Height of liquid depends on:

    1. Type of liquid (e.g., mercury, water).

    2. Atmospheric pressure.

Manometer

  • Device for measuring gas pressure in a container.

  • Compares gas pressure to atmospheric pressure:

    • Higher Height on Open End: Gas pressure is higher than atmospheric pressure.

    • Higher Height on Closed End: Gas pressure is lower than atmospheric pressure.

Gas Pressure Calculations

  • (a) Gas Pressure < Atmospheric Pressure:

    • Equation: Pgas = Patm - h

  • (b) Gas Pressure > Atmospheric Pressure:

    • Equation: Pgas = Patm + h

Additional Measurements of Pressure

  • Atmospheric Pressure: Varies with altitude and weather due to the mass of air being pulled toward the Earth by gravity.

Standard Atmospheric Pressure

  • SI unit for pressure: Pascals (Pa); commonly use kilopascals (kPa).

  • Standard values: 760.00 mmHg = 101.325 kPa = 1.0 atm.

  • Alternative: Pounds per square inch (lb/in² or psi).

Units of Pressure

  • Common Units:

    • mmHg: Used in manometers and barometers.

    • torr: Equal to mmHg.

    • standard atmosphere (atm).

    • Pascal (Pa).

    • KiloPascal (kPa).

  • Conversion: 1 atm = 760 mmHg = 760 torr = 101.3 kPa.

Example: Pressure Conversions

  • Given: Pressure of a gas = 49 torr.

    • Convert to atmospheres, Pascals, and mmHg.

    • Note: 1 atm = 760 mmHg = 760 torr = 101.3 kPa.

Kinetic Molecular Theory Recap

  • Review of initial kinetic molecular assumptions reiterated.

  • Importance of KMT in explaining gas laws.

Temperature and Kelvin Scale

  • Kelvin Temperature Scale: Extends Celsius scale down to absolute zero (0 K).

  • Absolute zero: Hypothetical point where no heat energy exists. Corresponds to -273.15°C.

  • Temperatures below absolute zero cannot exist (negative volumes).

  • Kelvin System: No negative values exist in this scale.

Kelvin Temperature Format

  • Temperature expressed in kelvins, not as degrees kelvin.

  • Conversion from Celsius to Kelvin: Add 273.

  • Baseline temperatures provided across temperature scales.