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) and entropy (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): Measure of molecular disorder. Melting increases entropy; freezing decreases entropy.

• 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 (ΔH > 0), increases disorder (ΔS > 0).

8.2 Intermolecular Forces

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

• London Dispersion Forces: Present in all molecules; strength increases with molecular size and surface area.

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

• Hydrogen Bonding: Strong dipole-dipole interaction between H bonded to O, N, or F and 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.

• Gas particles move randomly with negligible attractive forces.

• Volume of particles is much less than the space between them.

• Average kinetic energy is proportional to temperature (Kelvin).

• Collisions are elastic; total kinetic energy remains constant.

• An ideal gas obeys all these assumptions.

8.4 Pressure

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

• Atmospheric Pressure: Pressure exerted by air at sea level (1 atm = 760 mmHg = 101,325 Pa).

• Barometer: Measures atmospheric pressure using mercury column height.

• Manometer: Measures gas pressure in containers.

• SI Unit: Pascal (Pa).

Example: 55 psi = 3.74 atm = 379,212 Pa (conversion factors used).

8.5 Boyle’s Law: The Relation between Volume and Pressure

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

• If pressure increases, volume decreases proportionally.

• Mathematical expression:

• For two states:

Example: Compressing a gas from 1.0 atm to 9.5 atm reduces its volume from 750 mL to .

8.6 Charles’s Law: The Relation between 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.

• If temperature doubles, volume doubles.

• Mathematical expression:

• For two states:

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

8.7 Gay-Lussac’s Law: The Relation between 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.

• If temperature doubles, pressure doubles.

• Mathematical expression:

• For two states:

Example: Heating an aerosol can increases internal pressure.

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.

• Mathematical expression:

• Allows calculation of changes in pressure, volume, or temperature.

Example: Changing the volume and pressure of helium gas requires adjusting temperature to maintain conditions.

8.9 Avogadro’s Law: The Relation between Volume and Molar Amount

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

• Mathematical expression:

• Standard Temperature and Pressure (STP): 0°C (273.15 K), 1 atm.

• Standard molar volume at STP: 22.4 L per mole.

Example: Calculating moles of air in a room using its volume and the molar volume at STP.

8.10 The Ideal Gas Law

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

• Mathematical expression:

• R (gas constant): 0.0821 L·atm/(mol·K) or 8.314 J/(mol·K)

• Allows calculation of P, V, n, or T for an ideal gas.

Example: Finding moles of air in lungs or pressure in a gas cylinder using known values and the ideal gas law.

8.11 Partial Pressure and Dalton’s Law

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

• Mathematical expression:

• Each gas acts independently; its partial pressure depends on its concentration.

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

8.12 Liquids

Liquids are characterized by constant molecular motion and intermolecular forces. Properties like vapor pressure, boiling point, viscosity, and surface tension depend on these forces.

• 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 vs. interior.

Example: Water’s high surface tension and boiling point are due to hydrogen bonding.

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 (e.g., copper).

• 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 specific heats and heats of fusion/vaporization.

• 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.

• When not at phase change temperature, heat changes temperature; at phase change, heat overcomes intermolecular forces.

• Strong intermolecular forces mean higher heats of fusion and vaporization.

Example: To melt 0.300 mol naphthalene, convert moles to grams, then multiply by heat of fusion.

Concept Map

States of matter and their changes are interconnected through intermolecular forces, phase change temperatures, and energy calculations. Understanding these relationships is essential for predicting physical properties and behaviors of substances.