kinetic molecular theory NOTES

The Gas Laws

  • Introduction to Gas Laws, focusing on gas properties and their relationship with kinetic molecular theory.

Kinetic Molecular Theory of Gases

  • Definition: A theory explaining the properties of gases (pressure, temperature, volume) based on their composition and movement.

  • Key Concept: Kinetic refers to motion; the energy due to motion is known as kinetic energy (e.g., a rolling ball).

Components of Kinetic Theory

  • Three Main Components:

    1. Perfectly Elastic Collisions: No energy loss when gas molecules collide.

    2. No Volume: Gas molecules are extremely small, taking up negligible space.

    3. Constant Motion: Gas molecules are always moving in random, linear paths.

Gas Pressure

  • Definition: Gas pressure is caused by moving gas particles colliding with container walls.

  • Relation to Kinetic Theory: More collisions result in higher pressure.

Temperature and Kinetic Energy

  • Relation: Temperature measures the average kinetic energy of gas particles.

  • Higher Energy = Higher Temperature: As energy increases, temperature increases.

Factors Affecting Gas Pressure

  1. Number of Particles: Increasing gas particles leads to more collisions, increasing pressure.

    • Example: Doubling the number of gas particles doubles the pressure.

  2. Temperature: Higher temperatures increase kinetic energy, resulting in faster, more forceful collisions and higher pressure.

  3. Volume: Changes in volume affect pressure:

    • Increased Volume: Lowers pressure.

    • Decreased Volume: Increases pressure.

Barometers and Atmospheric Pressure

  • Evangelista Torricelli: Invented the barometer to measure atmospheric pressure (barometric pressure).

  • Measurement Unit: Pressure is measured in mm Hg using a column of mercury.

  • Standard Atmospheric Pressure: 1 atm = 760 mm Hg.

Pressure Measurement Units

  • SI Unit: Pascal (Pa)

  • Conversion: 1 atm = 101.3 kPa, 1 atm = 760 mm Hg, 1 atm = 14.7 psi.

Dimensional Analysis

  • Purpose: A method to change measurement units without changing the value, using conversion factors.

Pressure Conversion Factors

  • 1 atm = 760 mmHg = 14.7 psi = 101.325 kPa.

Examples of Dimensional Analysis

  • Converting 600 mm Hg to atm:

    • 600 mm x (1.00 atm/760 mm) = 0.789 atm.

  • Converting 65 psi to kPa:

    • 65 psi x (101.3 kPa/14.7 psi) = 447.93 kPa.

Understanding STP (Standard Temperature and Pressure)

  • Definition: Standard conditions at 1 atm and 273 K.

  • Importance: Essential for calculations involving gas laws.

Gas Laws Derived from Kinetic Molecular Theory

  • Key Laws:

    1. Boyle’s Law

    2. Charles’ Law

    3. Gay-Lussac’s Law

    4. Avogadro’s Law

    5. Ideal Gas Law

    6. Dalton’s Law

Ideal Gas Law (PV=nRT)

  • Variables:

    • P: Pressure (atm)

    • V: Volume (L)

    • n: Number of moles

    • T: Temperature (K)

    • R: Ideal gas constant (0.08206 L·atm/mol·K)

Examples of Ideal Gas Law Applications

  • Example 1: Calculate moles of oxygen in a defined volume, pressure, and temperature:

    • PV = nRT → n = 0.123 moles.

  • Example 2: Calculate volume occupied by 12.4 grams of O2 gas:

    • Use ideal gas law to find V.

Significance of STP in Problem Solving

  • Application of STP in calculations of gas behavior and volume adjustments under different conditions.

Dalton’s Law of Partial Pressures

  • Concept: Each gas in a mixture exerts pressure independently, called partial pressure.

  • Formula: Ptotal = P1 + P2 + P3 ...

  • Example Calculation: Total pressure of CO2, O2, and H2 mix:

    • Ptotal = 22.3 kPa + 44.7 kPa + 112 kPa = 179 kPa.

Gas Diffusion and Effusion

  • Diffusion: Mixing of gas molecules.

  • Effusion: Escape of gas through tiny holes; lighter gases effuse more rapidly.

  • Example: HCl and NH3 diffusion forming NH4Cl at different rates due to molar mass.