Characteristics of Gases

Overview of Gases

• Examples of Gases:

  • Monoatomic (Noble gases): He, Ne, Ar, Kr, Xe, Rn

  • Molecular (homonuclear diatomics): H2, O2, N2, F2, Cl2

  • Other gases: CH4, NO2, NO, N2O, CO2, NH3, HCl

  • Vapors (liquids with boiling points above room temperature): H2O, HNO3

Properties of Gases

• Volume: Gases occupy the full volume of their containers, filling any available space.

• Interdependence: Gas volume, pressure, and temperature are related; changing one affects the others.

• Density: Gases have low densities and viscosities compared to solids and liquids.

• Miscibility: Gases are highly miscible and can form homogeneous mixtures with other gases.

• Physical vs. Chemical Properties: Gases can have similar physical properties but very different chemical properties.

Gas Pressure and Its Measurement

Units of Pressure

• Pressure Equations: Pressure (P) = Force (F) / Area (A)

• Pressure Units:

  • 1 atm = 760 mm Hg = 760 torr = 1.013 × 10^5 Pa = 101.3 kPa = 1.013 bar

• Example Calculation:

  • Atmospheric pressure at 29.12 inches of mercury converts to 739.6 torr.

Measuring Pressure

• Barometers: Measure atmospheric pressure; invented by E. Torricelli in the 17th century.

• Manometers: Instrument for measuring gas pressure in a closed volume using a U-tube containing mercury.

  • Closed-end U-tube: Measures absolute pressure, Pgas = PΔh.

  • Open-end U-tube: Measures relative pressure, can indicate whether gas pressure is above or below atmospheric pressure.

The Gas Laws

Boyle’s Law

• Definition: At constant temperature, volume (V) is inversely proportional to pressure (P) for a given amount of gas: P1V1 = P2V2.

• Graphical Representation: Plotting V vs. P shows a hyperbolic relationship.

Charles’ Law

• Definition: At constant pressure, volume of a gas is directly proportional to its absolute temperature (T): V/T = constant.

• Key Point: Absolute zero is the temperature at which a gas would theoretically occupy zero volume.

Avogadro’s Law

• Definition: Volume of a gas is directly proportional to the number of moles (n) of the gas at constant temperature and pressure: V/n = constant.

• Implication: More gas means a larger volume under the same conditions.

The Ideal Gas Law

Gas Law Equation

• Combination of earlier laws leads to: PV = nRT, where R = ideal gas constant.

  • Common Values:

    • R = 0.0821 L·atm/(K·mol)

    • R = 62.36 L·torr/(K·mol)

• Standard Conditions: At STP (0℃ and 1 atm), 1 mole of an ideal gas occupies 22.41 L.

Applications and Examples

• Gas Density and Molar Mass: Relates density to molar mass using equations involving P, V, T, and R.

Kinetic-Molecular Theory

Fundamental Postulates

• Gases consist of numerous particles in random motion.

• Gas particles have negligible volume and negligible intermolecular forces.

• Collisions between gas particles and with container walls are elastic.

Explanation of Gas Laws

• Pressure Variation: Pressure depends on collision frequency and particle speed; increasing temperature raises pressure due to increased particle motion.

Effusion and Diffusion

Definitions

• Effusion: Escape of gas particles through orifices.

• Diffusion: Spreading of gas particles in space due to thermal motion.

Graham’s Law of Effusion

• States that the rate of effusion is inversely proportional to the square root of the gas's molar mass.

Non-Ideal Gas Behavior

Real Gases

• Deviate from ideal behavior due to molecular volume and intermolecular forces, particularly at high pressures and low temperatures.

van der Waals Equation

• Adjusts the ideal gas law to account for non-ideal behavior with parameters specific to the particular gas.