CHEM 11_Chapter 11 - Properties of Gases

11.1 Gases: Properties and Behavior

  • Key Terms

    • Kinetic Molecular Theory of Gases:

      • Scientific model explaining the behavior of gases.

      • Based on assumptions about particle motion, energy, and interactions.

    • Ideal Gas:

      • Hypothetical gas defined by specific characteristics:

        • High translational kinetic energy.

        • No attractive or repulsive forces between particles.

        • Collisions are elastic, conserving total kinetic energy.

        • Particle volume is negligible compared to the container's volume.

        • Average kinetic energy is directly related to temperature.

    States of Matter and Particle Behavior

    • Solids:

      • Constant shape and volume, almost incompressible.

      • Particles:

        • Arranged in a fixed, regular pattern.

        • Vibrate in place but do not move past each other.

    • Liquids:

      • Variable shape, constant volume, almost incompressible.

      • Particles:

        • Less organized, able to slide past each other.

        • Allow liquids to flow and take the shape of their container.

    • Gases:

      • Variable shape and volume, compressible.

      • Particles:

        • Highly disorganized, widely separated, move freely in all directions.

        • Bounce off each other and the container walls.

    Changes of State

    • Determined by two factors:

      1. Attractive Forces: Hold particles together.

        • Stronger forces require more kinetic energy to overcome.

      2. Kinetic Energy: Tends to pull particles apart.

    Attractive Forces That Influence the State of Matter

    Stronger:

    • Attractions between oppositely charged particles (IONS)

      • State: Solid

      • Example: Table salt (NaCl(s))

    • Attractions between polar molecules

      • State: Solid, liquid, gas

      • Examples:

        • Glucose (C6H12O6(s))

        • Ethanol (CH3CH2OH())

        • Ammonia (NH3(g))

    Weaker:

    • Attractions between non-polar molecules

      • State: Solid, liquid, gas

      • Examples:

        • Paraffin (C30H62(s))

        • Pentane (C5H12())

        • Carbon dioxide (CO2(g))

    Attractive Forces Between Particles

    1. Attractive Forces Between Oppositely Charged Particles:

      • Electrostatic attractions: Oppositely charged particles pull towards each other (like magnets).

      • Example: Ionic compounds have very strong bonds due to these attractions. A positive ion is attracted to the negative ion.

      • Result: Ionic compounds, like NaCl, are solid at room temperature.

    2. Attractive Forces Between Polar Molecules:

      • Polar molecules have partial positive and partial negative ends, creating a permanent dipole-dipole effect.

      • Dipole bonds between polar molecules are weaker than ionic bonds.

      • Result: Substances with polar molecules often remain as liquids or gases at room temperature.

      • Hydrogen Bonds:

        • A type of dipole-dipole interaction.

        • Strength depends on the type of bond:

          • O–H bonds are stronger than N–H bonds.

        • Example: Water (liquid at room temperature) has stronger O–H hydrogen bonds than ammonia (gas at room temperature).

    3. Attractive Forces Between Non-Polar Molecules:

      • Forces arise from temporary dipoles.

      • Large non-polar molecules tend to be liquids at room temperature because their size increases attractive forces.

      • Key Relationships:

        • Larger molecules = Stronger attractive forces.

        • Stronger forces = More energy (heat) needed to cause a state change.

    Kinetic Energy and Motion of Particles

    • Three types of particle motion:

      1. Vibrational Motion: Back-and-forth movement (all states).

      2. Rotational Motion: Spinning on an axis (liquids, gases, some solids).

      3. Translational Motion: Movement from one place to another (liquids, gases).

    • Temperature:

      • Directly related to the average kinetic energy of particles.

      • Heating increases particle motion (vibrational, rotational, and translational).

      • Higher temperatures allow particles to overcome attractive forces, leading to state changes.

    Distinguishing Properties of Gases

    1. Compressibility:

      • Gases can be compressed significantly due to large spaces between particles.

      • Solids and liquids are nearly incompressible.

    2. Thermal Expansion:

      • Gases expand greatly with temperature increase (at constant pressure).

      • Solids and liquids expand minimally.

    3. Low Viscosity:

      • Gases flow freely due to minimal intermolecular forces.

    4. Low Density:

      • Gas densities are much lower than those of solids and liquids.

    5. Miscibility:

      • Gases mix completely with other gases due to negligible particle volume and high mobility.

    The Kinetic Molecular Theory of Gases

    1. Constant Random Motion:

      • Particles travel in straight lines until collisions occur.

    2. Negligible Volume:

      • Individual gas particles are considered point masses with negligible size.

    3. No Interparticle Forces:

      • No attractions or repulsions between particles.

    4. Elastic Collisions:

      • Total kinetic energy is conserved during collisions.

    5. Temperature-Kinetic Energy Relationship:

      • Average kinetic energy increases with temperature.

    Behavior of Real Gases vs. Ideal Gases

    • Ideal Gas Behavior: Accurate under ordinary temperatures and pressures.

    • Deviations: Occur at very high pressures or low temperatures where particle volume and intermolecular forces cannot be ignored.

11.2 Gases and Pressure Changes

Key Terms

  • Viscosity:

    • Resistance of a liquid to flow.

    • Influenced by temperature and intermolecular forces.

    • Example: Honey has higher viscosity than water.

  • Surface Tension:

    • The energy required to increase a liquid’s surface area.

    • Results from cohesive forces between molecules.

    • Example: Water droplets form spheres due to high surface tension.

  • Capillary Action:

    • The ability of a liquid to flow against gravity in a narrow tube.

    • Caused by the combination of adhesive forces (to the tube) and cohesive forces (within the liquid).

    • Example: Water rising in a thin straw.

Distinguishing Properties of Liquids

  1. Definite Volume, Indefinite Shape:

    • Liquids retain volume but take the shape of their container.

  2. Incompressibility:

    • Liquids resist compression due to closely packed particles.

  3. Fluidity:

    • Liquids flow due to their ability to slide past one another.

    • Lower viscosity = higher fluidity.

  4. Density:

    • Liquids are denser than gases but less dense than solids (in most cases).

Intermolecular Forces in Liquids

  1. Dipole-Dipole Interactions:

    • Found in polar molecules.

    • Moderate strength.

    • Example: Hydrogen bonds in water.

  2. London Dispersion Forces:

    • Present in all molecules, strongest in large, non-polar molecules.

    • Example: Dispersion forces in bromine (Br2).

  3. Hydrogen Bonding:

    • A special type of dipole interaction involving H bonded to N, O, or F.

    • Example: Water (H2O) has strong hydrogen bonds.

Factors Affecting Liquid Properties

  1. Temperature:

    • Higher temperatures reduce viscosity and surface tension.

    • Example: Heating honey makes it flow more easily.

  2. Intermolecular Forces:

    • Stronger forces lead to higher viscosity and surface tension.

    • Example: Water’s high surface tension due to hydrogen bonding.

  3. Particle Size and Shape:

    • Larger or irregularly shaped particles increase viscosity.

Capillary Action Explained

  • Adhesive Forces:

    • Attraction between liquid molecules and a solid surface.

    • Stronger adhesive forces cause liquids to "climb" surfaces.

  • Cohesive Forces:

    • Attraction between liquid molecules themselves.

    • Cohesion prevents the liquid from spreading too thin.

  • Examples:

    • Water rises in plant roots due to capillary action.

    • Mercury forms a convex meniscus due to stronger cohesion than adhesion.

Evaporation and Vapour Pressure

  1. Evaporation:

    • Surface particles with sufficient kinetic energy escape into the gas phase.

    • Increases with temperature and surface area.

  2. Vapour Pressure:

    • Pressure exerted by a liquid’s vapor in a closed system.

    • Higher temperatures = higher vapor pressure.

  3. Boiling Point:

    • The temperature at which vapor pressure equals external pressure.

    • Stronger intermolecular forces = higher boiling point.

Dynamic Equilibrium in Liquids

  • Definition:

    • When the rate of evaporation equals the rate of condensation in a closed system.

  • Example:

    • A sealed bottle of water reaches equilibrium between liquid and vapor phases.

11.3 Gases and Temperature Changes

  • Charles's Law

    • Direct Proportionality: Volume of a gas is directly proportional to its Kelvin temperature when pressure and the amount of gas (number of moles) remain constant.

      • This means that if the temperature of a gas increases, its volume will also increase proportionally, and vice versa, as long as pressure and the amount of gas stay the same.

    • Mathematical Expression:

      • V ∝ T (V is directly proportional to T)

      • V1/T1 = V2/T2 (where V1 and T1 are the initial volume and temperature, and V2 and T2 are the final volume and temperature)

    • Temperature in Kelvin:

      • Temperature must always be expressed in Kelvin (K) for Charles's Law to hold true.

      • K = °C + 273.15

    • Graph of V vs. T (in Kelvin):

      • A straight line passing through the origin (0,0) on a graph of volume (V) versus Kelvin temperature (T). This visually represents the direct proportionality.

    • Kinetic Molecular Theory Explanation:

      • Increased Temperature:

        • Leads to increased kinetic energy of the gas molecules.

        • Molecules move faster and collide more frequently and with greater force.

      • Constant Pressure:

        • To maintain constant pressure, the volume of the container must increase to accommodate the increased molecular motion.

        • This is because the increased collisions exert a greater force on the container walls, and a larger volume reduces the frequency and force of these collisions.

    Gay-Lussac's Law

    • Direct Proportionality: Pressure of a gas is directly proportional to its Kelvin temperature when volume and the amount of gas (number of moles) remain constant.

      • This means that if the temperature of a gas increases, its pressure will also increase proportionally, and vice versa, as long as volume and the amount of gas stay the same.

    • Mathematical Expression:

      • P ∝ T (P is directly proportional to T)

    • Temperature in Kelvin:

      • Temperature must always be expressed in Kelvin (K) for Gay-Lussac's Law to hold true.

    • Kinetic Molecular Theory Explanation:

      • Increased Temperature:

        • Leads to increased kinetic energy of the gas molecules.

        • Molecules move faster and collide more frequently and with greater force.

      • Constant Volume:

        • Since the volume of the container cannot change, the increased collisions result in a significant increase in pressure exerted on the container walls.

    Key Points for Both Laws

    • Ideal Gas Behavior: These laws describe the behavior of ideal gases. Real gases may exhibit slight deviations, particularly at high pressures or low temperatures.

    • Practical Applications: Both Charles's Law and Gay-Lussac's Law have significant practical applications in various fields, including:

      • Weather forecasting

      • Aviation

      • Industrial processes (e.g., manufacturing, chemical engineering)