Chapter 7: Gases, Liquids, and Solids – Study Notes

States of Matter

The Three States of Matter

Matter exists in three primary states: gas, liquid, and solid. Each state is characterized by distinct physical properties, particle arrangements, and types of movement.

• Gas: Expands to fill its container, particles are far apart and move rapidly, low density, negligible interactions between particles.

• Liquid: Has a fixed volume but takes the shape of its container, particles are close but can move past each other, higher density than gases, moderate interactions.

• Solid: Has a definite shape and volume, particles are closely packed in a fixed arrangement, high density, strong interactions.

Property	Gas	Liquid	Solid
Shape & Volume	Expands to fill container	Fixed volume, shape of container	Definite shape & volume
Arrangement	Random, far apart	Random, close	Fixed, very close
Density	Low (<0.01 g/mL)	High (~1 g/mL)	High (1–10 g/mL)
Particle Movement	Very fast	Moderate	Slow
Interactions	None	Strong	Very strong

Key Point: The state of a substance depends on the balance between the kinetic energy of its particles and the strength of intermolecular forces.

Gas Laws

Kinetic-Molecular Theory of Gases

The kinetic-molecular theory explains the behavior of gases:

• Gas particles move randomly and rapidly.

• They are small compared to the space between them.

• No attractive forces exist between particles.

• Kinetic energy increases with temperature.

• Collisions are elastic; particles rebound without energy loss.

Gas Pressure

Gas pressure is the force exerted by gas particles colliding with the walls of a container. It is defined as:

$P = \frac{F}{A}$

• 1 atmosphere (atm) = 760 mm Hg = 760 torr = 14.70 psi = 101,325 Pa

Boyle’s Law

For a fixed amount of gas at constant temperature, pressure and volume are inversely related:

$P \times V = k$

or for two sets of conditions:

$P_1 V_1 = P_2 V_2$

Example: If the volume of a gas is halved, the pressure doubles.

Application: Breathing

Inhalation increases lung volume, decreasing pressure and drawing air in. Exhalation decreases volume, increasing pressure and expelling air.

Charles’s Law

For a fixed amount of gas at constant pressure, volume is proportional to Kelvin temperature:

$\frac{V}{T} = k$

or

$\frac{V_1}{T_1} = \frac{V_2}{T_2}$

Example: Doubling the temperature (in K) doubles the volume.

Gay-Lussac’s Law

For a fixed amount of gas at constant volume, pressure is proportional to Kelvin temperature:

$\frac{P}{T} = k$

or

$\frac{P_1}{T_1} = \frac{P_2}{T_2}$

Combined Gas Law

Combines Boyle’s, Charles’s, and Gay-Lussac’s laws:

$\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

Avogadro’s Law

At constant temperature and pressure, the volume of a gas is proportional to the number of moles:

$\frac{V}{n} = k$

or

$\frac{V_1}{n_1} = \frac{V_2}{n_2}$

At STP (0°C, 1 atm), 1 mole of any gas occupies 22.4 L (standard molar volume).

The Ideal Gas Law

Relates pressure, volume, temperature, and moles:

$PV = nRT$

• R = 0.08206 L·atm/(mol·K) or 62.37 L·mmHg/(mol·K)

Dalton’s Law of Partial Pressures

The total pressure of a gas mixture is the sum of the partial pressures of each component:

$P_{\text{total}} = P_A + P_B + P_C + \ldots$

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules, influencing physical properties like boiling and melting points.

• London Dispersion Forces: Weak, present in all molecules due to temporary dipoles.

• Dipole-Dipole Interactions: Moderate, occur between polar molecules with permanent dipoles.

• Hydrogen Bonding: Strongest, occurs when H is bonded to O, N, or F and attracted to O, N, or F in another molecule.

Type of Force	Relative Strength	Exhibited by	Example
London dispersion	Weak	All molecules	CH4, H2CO, H2O
Dipole-dipole	Moderate	Polar molecules	H2CO, H2O
Hydrogen bonding	Strong	O–H, N–H, or H–F bonds	H2O

Boiling and Melting Points

The stronger the intermolecular forces, the higher the boiling and melting points of a substance.

Liquids: Properties and Behavior

Evaporation and Condensation

Evaporation is the endothermic process where a liquid becomes a gas. Condensation is the exothermic process where a gas becomes a liquid.

Vapor Pressure

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid. It increases with temperature and is lower for substances with strong intermolecular forces.

Viscosity and Surface Tension

Viscosity is a liquid's resistance to flow; higher with stronger intermolecular forces or larger molecules. Surface tension is the resistance of a liquid to spread out, also higher with stronger forces.

Solids: Structure and Types

Crystalline and Amorphous Solids

Solids can be crystalline (ordered, repeating structure) or amorphous (no regular arrangement). Crystalline solids are further classified as ionic, molecular, network, or metallic.

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

• Molecular: Composed of molecules (e.g., H2O).

• Network: Atoms covalently bonded in a network (e.g., SiO2).

• Metallic: Metal cations in a sea of electrons (e.g., Cu).

Energy and Phase Changes

Specific Heat

Specific heat is the energy required to raise the temperature of 1 g of a substance by 1°C:

$\text{heat} = \text{mass} \times \Delta T \times \text{specific heat}$

Phase Changes and Heats of Transformation

• Heat of fusion: Energy to melt 1 g of a solid.

• Heat of vaporization: Energy to vaporize 1 g of a liquid.

Process	Phases	Energy Change
Melting	Solid → liquid	Endothermic (absorbs energy)
Freezing	Liquid → solid	Exothermic (releases energy)
Vaporization	Liquid → gas	Endothermic
Condensation	Gas → liquid	Exothermic
Sublimation	Solid → gas	Endothermic
Deposition	Gas → solid	Exothermic

Heating and Cooling Curves

A heating curve shows temperature changes as heat is added; a cooling curve shows temperature changes as heat is removed. Plateaus indicate phase changes where temperature remains constant as energy is used to change state.