Chapter 8: Gases, Liquids, and Solids

8.1 States of Matter and Their Changes

The three primary states of matter—solid, liquid, and gas—are distinguished by the arrangement and energy of their particles. Changes between these states are called phase changes and involve energy transfer as heat.

• Change of State: The transformation of a substance from one state (solid, liquid, gas) to another.

• Enthalpy (ΔH): The heat absorbed or released during a phase change.

• Entropy (ΔS): The change in molecular disorder or freedom during a process.

• Melting (Fusion): Solid to liquid; heat is absorbed (endothermic, ΔH > 0), and disorder increases (ΔS > 0).

• Freezing: Liquid to solid; heat is released (exothermic, ΔH < 0), and disorder decreases (ΔS < 0).

• Equilibrium Temperature: The temperature at which the enthalpy and entropy changes balance, resulting in equilibrium between phases (e.g., melting point, boiling point).

Example: Ice melting at 0°C absorbs heat and increases molecular disorder.

8.2 Intermolecular Forces

Intermolecular forces (IMFs) are attractions between molecules or atoms that influence physical properties such as boiling and melting points.

• London Dispersion Forces: Present in all molecules; arise from temporary shifts in electron density. Strength increases with molecular weight and surface area.

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

• Hydrogen Bonding: A strong dipole-dipole interaction between hydrogen bonded to O, N, or F and another O, N, or F atom.

Physical Properties: Stronger IMFs lead to higher boiling and melting points.

Example: Methane (nonpolar) has only London forces; HCl (polar) has dipole-dipole and London forces; acetic acid (polar, with O-H) has all three, including hydrogen bonding. Boiling points increase in this order.

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 and rapidly with negligible attractive forces.

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

• Average kinetic energy is proportional to temperature (in kelvins).

• Collisions are elastic (no energy lost).

• An ideal gas perfectly follows these assumptions.

8.4 Pressure

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

• Atmospheric Pressure: The weight of air above Earth's surface; standard is 1 atm = 760 mmHg = 101,325 Pa.

• Barometer: Measures atmospheric pressure using mercury column height.

• Manometer: Measures gas pressure in a container relative to atmospheric pressure.

• SI Unit: Pascal (Pa); 1 atm = 101,325 Pa.

Example: To convert 55 psi to atm:

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

Boyle’s Law states that for a fixed amount of gas at constant temperature, the volume is inversely proportional to pressure.

• Mathematical form:

• For two states:

Example: Compressing a gas from 1.0 atm to 9.5 atm reduces its volume by a factor of 9.5.

8.6 Charles’s Law: The Relation between Volume and Temperature

Charles’s Law states that for a fixed amount of gas at constant pressure, the volume is directly proportional to the absolute temperature (in kelvins).

• Mathematical form:

• For two states:

Example: Heating air in a balloon causes it to expand as temperature increases.

8.7 Gay-Lussac’s Law: The Relation between Pressure and Temperature

Gay-Lussac’s Law states that for a fixed amount of gas at constant volume, the pressure is directly proportional to the absolute temperature.

• Mathematical form:

• For two states:

Example: Heating a sealed aerosol can increases the 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 form:

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

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

Avogadro’s Law states that at constant temperature and pressure, the volume of a gas is directly proportional to the number of moles (n).

• Mathematical form:

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

• Standard molar volume at STP: 22.4 L/mol for an ideal gas.

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 one equation.

• Equation:

• R (gas constant): or

Example: Finding the number of moles of air in the lungs at a given volume, pressure, and temperature.

8.11 Partial Pressure and Dalton’s Law

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

• Equation:

• Partial pressure can be calculated as: where is the mole fraction.

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

8.12 Liquids

Liquids have particles in constant motion, with intermolecular forces holding them together but allowing flow.

• Vapor Pressure: The pressure exerted by vapor in equilibrium with its liquid.

• Vapor pressure increases with temperature and depends on the strength of IMFs.

• Boiling Point: The temperature at which vapor pressure equals atmospheric pressure.

• Viscosity: Resistance to flow; increases with stronger IMFs.

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

8.13 Solids

Solids are characterized by fixed positions of their particles. They can be crystalline or amorphous.

• Crystalline Solids: Ordered, repeating structures. Types include:

  • Ionic Solids: Composed of ions (e.g., NaCl).

  • Molecular Solids: Composed of molecules held by IMFs (e.g., ice).

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

  • Metallic Solids: Metal cations in a sea of electrons; conduct electricity, malleable.

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

8.14 Changes of State Calculations

Calculating the heat required for phase changes involves the heat of fusion (melting) and heat of vaporization (boiling).

• Heat of Fusion (ΔHfus): Heat required to melt 1 g (or 1 mol) of a substance at its melting point.

• Heat of Vaporization (ΔHvap): Heat required to vaporize 1 g (or 1 mol) of a substance at its boiling point.

• When not at a phase change temperature, heat changes temperature:

• At phase change, heat is used to overcome IMFs: or

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

Concept Map

A concept map for this chapter would connect the states of matter, phase changes, intermolecular forces, and the gas laws, showing how each property and law relates to the behavior of substances in different states.