Focus Guide for Exam IV CH 10 & 11

Focus Guide for EXAM IV: CH 10 & 11

General Characteristics about Gases

  • Gases have no definite shape or volume; they expand to fill their container.

  • The density of gases is significantly lower than that of solids and liquids.

  • Gases are highly compressible due to the large amount of empty space between molecules.

  • Gas particles are in constant motion, colliding with each other and the walls of their container.

  • Temperature, pressure, and volume significantly influence the behavior of gases.

Kinetic Molecular Theory (KMT) Parts

  • Postulate 1: Gases consist of large numbers of tiny particles that are far apart relative to their size.

  • Postulate 2: Gas particles are in continuous, rapid, and random motion.

  • Postulate 3: Collisions between gas particles and between particles and the container walls are perfectly elastic, meaning kinetic energy is conserved.

  • Postulate 4: The average kinetic energy of gas particles is directly proportional to the gas temperature in Kelvin.

  • Postulate 5: There are no attractive or repulsive forces between gas particles; they interact only during collisions.

Conversion of Pressure Units

  • Understand the different pressure units:

    • Atmospheres (atm)

    • Pascals (Pa)

    • Millimeters of mercury (mmHg or torr)

    • Psi (pounds per square inch)

  • Conversion Equivalents:

    • 1 atm = 101.325 kPa = 760 mmHg = 760 torr = 14.696 psi

Determining Identity of a Gas

  • Given the mass (grams), pressure (atm or mmHg), temperature (Kelvin), and volume (liters), use the Ideal Gas Law:

  • PV = nRT

    • Where:

    • P = pressure

    • V = volume

    • n = number of moles

    • R = ideal gas constant (0.0821 L·atm/(K·mol))

    • T = temperature in Kelvin

  • Calculate number of moles (n) to identify the gas.

Graham’s Law of Effusion Formula

  • Formula: \frac{Rate\,of\,effusion\,of\,gas1}{Rate\,of\,effusion\,of\,gas2} = \frac{\sqrt{M2}}{\sqrt{M1}}

    • Where M is the molar mass of the gas.

  • This law states that lighter gases effuse faster than heavier gases.

Ideal Gas Law Formula and Usage

  • The Ideal Gas Law: PV = nRT

    • Key to solving problems involving gases under various conditions.

  • Use this law to find pressure, volume, temperature, or number of moles when the other variables are known.

Combined Gas Law Formula and Usage

  • Combination of Boyle’s Law, Charles’s Law, and Gay-Lussac’s Law:

  • Formula: \frac{P1V1}{T1} = \frac{P2V2}{T2}

  • This allows calculations when only two states of a gas are considered.

Dalton’s Law of Partial Pressure and Usage

  • Definition: The total pressure of a gas mixture is equal to the sum of the partial pressures of individual gases.

  • Formula: P{total} = P1 + P2 + P3 + …

  • Useful in calculating the contributions of various gases in a mixture.

Calculating Mole Fraction from Partial Pressure

  • Mole Fraction: Xi = \frac{Pi}{P_{total}}

  • Where $Xi$ is the mole fraction of gas i and $Pi$ is the partial pressure of gas i.

  • Important for determining the composition of gas mixtures.

Solving Gas Collected Over Water Problem

  • When gases are collected over water, total pressure is the sum of the water vapor pressure and the pressure of the gas:

  • P{total} = P{gas} + P{H2O}

  • Adjust calculations by subtracting the water vapor pressure from total pressure to find the pressure of the gas.

Conditions for Real Gases to Behave Like Ideal Gases

  • Real gases behave most like ideal gases under:

    • High temperature

    • Low pressure

  • Under these conditions, intermolecular forces and the volume of gas particles become negligible.

Intermolecular Forces Differences

  • Hydrogen Bonding: Strong attraction between molecules that contain a hydrogen atom bonded to a highly electronegative atom (N, O, or F).

  • Dipole-Dipole Interactions: Attractions between polar molecules due to their positive and negative ends.

  • London Dispersion Forces (LDFs): Weak intermolecular forces arising from the temporary dipoles in nonpolar molecules.

  • Ion-Dipole Forces: Attractive forces between an ion and a polar molecule; significant in solutions.

  • Ionic Forces: Strong attractions between positively and negatively charged ions.

Reading a Phase Diagram

  • Triple Point: The unique condition at which all three phases (solid, liquid, gas) coexist in equilibrium.

  • Critical Point: The endpoint of a phase equilibrium curve, above which the substance cannot exist as a liquid.

  • Understand where solids, liquids, and gases are represented in the phase diagram.

Reading a Heating/Cooling Curve

  • A graph showing the phase changes of a substance over time as heat is added or removed.

  • Key points:

    • Melting and freezing points

    • Boiling and condensation points

    • Plateau regions indicate phase changes without temperature change.

Exothermic vs. Endothermic Processes

  • Exothermic: Processes that release heat to the surroundings (e.g., combustion).

  • Endothermic: Processes that absorb heat from the surroundings (e.g., photosynthesis).

General Characteristics of Solids, Liquids, and Gases

  • Solids: Have a definite shape and volume, particles are closely packed and vibrate in place.

  • Liquids: Have a definite volume but take the shape of their container, particles are closely packed but can move past one another.

  • Gases: No definite shape or volume, particles are far apart and move freely.

Definitions of Key Terms

  • Viscosity: A measure of a fluid's resistance to flow. High viscosity indicates a thick fluid, while low viscosity indicates a thin fluid.

  • Vapor Pressure: The pressure exerted by a vapor in equilibrium with its liquid at a given temperature.

  • Surface Tension: The energy required to increase the surface area of a liquid due to intermolecular forces.