Gases, Liquids, and Solids

8.1 States of Matter and Their Changes

Matter exists in three primary states: solid, liquid, and gas. The state of a substance depends on the balance between the attractive forces among particles and their kinetic energy. Changes of state (phase changes) are reversible and involve changes in enthalpy (heat content) and entropy (molecular disorder).

• Phase Changes: Transitions between solid, liquid, and gas (e.g., melting, freezing, boiling).

• Enthalpy (ΔH): Heat absorbed or released during a phase change. Endothermic changes (e.g., melting) absorb heat; exothermic changes (e.g., freezing) release heat.

• Entropy (ΔS): Change in molecular disorder. Melting increases disorder (ΔS positive); freezing decreases disorder (ΔS negative).

• Melting Point: Temperature at which solid and liquid phases are in equilibrium.

• Boiling Point: Temperature at which liquid and gas phases are in equilibrium.

Example: Melting ice absorbs heat (endothermic) and increases disorder (entropy).

8.2 Intermolecular Forces

Intermolecular forces are attractions between molecules or atoms, influencing physical properties like melting and boiling points. They are negligible in gases but significant in liquids and solids.

• London Dispersion Forces: Present in all molecules; arise from temporary polarization. Stronger in larger molecules.

• Dipole-Dipole Forces: Occur in polar molecules; positive and negative ends attract.

• Hydrogen Bonding: Strong dipole-dipole interaction; occurs when H is bonded to O, N, or F and interacts with another O, N, or F atom.

Example: Water forms four hydrogen bonds per molecule, leading to high boiling point.

8.3 Gases and the Kinetic-Molecular Theory

The kinetic-molecular theory explains gas behavior based on particle motion and energy.

• Assumptions: Gas particles move randomly, have negligible volume, and no attractive forces.

• Elastic Collisions: Total kinetic energy remains constant.

• Temperature: Average kinetic energy is proportional to Kelvin temperature.

• Ideal Gas: A gas that perfectly follows these assumptions.

8.4 Pressure

Pressure is the force per unit area exerted by gas particles colliding with surfaces.

• Atmospheric Pressure: Pressure exerted by air; measured with a barometer.

• Units: Common units include mmHg (torr), atm, and pascal (Pa).

• Manometer: Measures gas pressure in a container.

Conversion Example: 1 atm = 760 mmHg = 101,325 Pa.

8.5 Boyle’s Law: Volume and Pressure

Boyle’s law describes the inverse relationship between volume and pressure for a fixed amount of gas at constant temperature.

• Equation:

• Alternate Form:

• Application: Compressing a gas increases its pressure and decreases its volume.

Example: Compressing air in a car engine from 1.0 atm to 9.5 atm reduces volume proportionally.

8.6 Charles’s Law: Volume and Temperature

Charles’s law states that the volume of a gas is directly proportional to its Kelvin temperature at constant pressure and amount.

• Equation:

• Alternate Form:

• Application: Heating a gas increases its volume.

Example: Air inhaled at room temperature expands when warmed to body temperature.

8.7 Gay-Lussac’s Law: Pressure and Temperature

Gay-Lussac’s law states that the pressure of a gas is directly proportional to its Kelvin temperature at constant volume and amount.

• Equation:

• Alternate Form:

• Application: Heating a sealed gas container increases its pressure.

Example: Aerosol can pressure rises when heated.

8.8 The Combined Gas Law

The combined gas law merges Boyle’s, Charles’s, and Gay-Lussac’s laws for a fixed amount of gas.

• Equation:

• Application: Used when pressure, volume, and temperature all change.

Example: Calculating the temperature needed to maintain a certain pressure and volume.

8.9 Avogadro’s Law: Volume and Moles

Avogadro’s law states that the volume of a gas is directly proportional to the number of moles at constant pressure and temperature.

• Equation:

• Alternate Form:

• Standard Molar Volume: At STP (0°C, 1 atm), 1 mole of any ideal gas occupies 22.4 L.

Example: Calculating moles of air in a room using its volume.

8.10 The Ideal Gas Law

The ideal gas law combines all gas variables into a single equation.

• Equation:

• Variables: P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K).

• Gas Constant (R): or

Example: Calculating moles of air in lungs or pressure in a gas cylinder.

8.11 Partial Pressure and Dalton’s Law

Partial pressure is the pressure a single gas in a mixture would exert if it were alone. Dalton’s law states that the total pressure is the sum of the partial pressures of all gases present.

• Equation:

• Application: Used to calculate partial pressures in air or gas mixtures.

Example: Partial pressures of oxygen, nitrogen, water vapor, and argon in humid air.

8.12 Liquids

Liquids have molecules in constant motion. Some escape into the gas phase (vapor), and in a closed container, evaporation and condensation reach equilibrium.

• Vapor Pressure: Partial pressure of vapor in equilibrium with liquid; increases with temperature.

• Boiling Point: Occurs when vapor pressure equals atmospheric pressure.

• Viscosity: Resistance to flow; increases with stronger intermolecular forces.

• Surface Tension: Caused by differences in intermolecular forces at the surface versus interior.

8.13 Solids

Solids are classified as crystalline or amorphous based on particle arrangement.

• Crystalline Solids: Ordered arrangement; subdivided into:

  • Ionic Solids: Alternating ions held by ionic bonds (e.g., NaCl).

  • Molecular Solids: Molecules held by intermolecular forces (e.g., ice).

  • Covalent Network Solids: Atoms linked by covalent bonds in a giant array (e.g., diamond).

  • Metallic Solids: Metal cations in a sea of electrons; malleable and conductive.

• Amorphous Solids: Random arrangement; no long-range order (e.g., glass, tar).

8.14 Changes of State Calculations

Calculating heat associated with phase changes involves the heat of fusion (melting) and heat of vaporization (boiling).

• Heat of Fusion: Heat required to melt 1 g of a substance at its melting point.

• Heat of Vaporization: Heat required to vaporize 1 g of a substance at its boiling point.

• Calculation: Total heat = mass × heat of fusion or vaporization.

• Strong Intermolecular Forces: Lead to higher heats of fusion and vaporization.

Example: Melting naphthalene: Convert moles to grams, then multiply by heat of fusion to find total heat required.

Concept Map

A concept map links the relationships among states of matter, intermolecular forces, gas laws, and phase changes, showing how each topic connects to others in the chapter.