Chapter 12

Matter exists in three primary phases: solid, liquid, and gas.
Physical properties of each phase are influenced by the arrangement and interaction of particles.

Table: States of Matter Comparison
- Gas: Conforms to shape and volume, high compressibility, flows easily.
- Liquid: Conforms to the shape of the container, fixed volume, low compressibility, flows moderately.
- Solid: Fixed shape and volume, negligible compressibility, does not flow.

Defined as a physically distinct, homogeneous part of a system.
Properties determined by the balance between potential and kinetic energy of particles.
- Potential Energy: Attractive forces that pull particles together.
- Kinetic Energy: Motion that tends to disperse particles.

Gas
- Weak attractive forces relative to kinetic energy. No fixed shape or volume.
Liquid
- Stronger attractive forces because particles have lower kinetic energy. Fixed volume, can flow and change shape.
Solid
- Attractions dominate motion; fixed shape and volume due to closely packed particles.

Intramolecular Forces: Forces within a molecule that determine chemical behavior.
- Example: Water (H2O) behaves similarly regardless of phase.
- bonding forces
Intermolecular Forces: Forces between molecules that determine physical behavior.
- Varies between different phases of matter based on strength.
- referred to as IMF or nonbonding forces

Types of Phase Changes
- Melting: Solid to liquid (endothermic)
- Freezing: Liquid to solid (exothermic)
- Vaporizing: Liquid to gas (endothermic)
- Condensing: Gas to liquid (exothermic)
- Sublimation: Solid to gas (endothermic)
- Deposition: Gas to solid (exothermic)
Enthalpy change indicates energy absorbed/released during phase changes.
- ΔH: Represents energy required to break or form interactions.

Intermolecular forces arise from attractions between partial charges in molecules (e.g., dipoles).
F = k ((n+e-)(n-e-)/d²)
- F = force of attractionn+ or
- n- = charge of cation or anion
- e- = charge of electron (1.602 x 10-19 Coulombs)
- d = distance between charges
- k = a constant
Ion-Dipole Forces: Attraction between ions and polar molecules.
Dipole-Dipole Forces: Attraction between polar molecules.
Hydrogen Bonds: Strong dipole-dipole attraction between H and electronegative atoms (N, O, F).
INCREASED charge = INCREASED force
DECREASED distance between charges = INCREASED force
Magnitude of charges: ION > DIPOLE > “INDUCED” DIPOLE
- Ion = change in e- , uneven distribution of charge
  - Whole number charge (+1, -3)
  - loss/gain of e-
- Dipole = change in distribution of e- cloud (orbitals), even distribution of charge
  - Partial charges (< +- 1)
  - two charges/ends
  - polar covalent
- induced dipole = change in distribution of e- cloud (orbitals)
  - Neutral
  - weak
  - non-polar covalent
Increasing strength;
- Ion + dipole (strongest)
- Dipole + Dipole
- Dipole + induced dipole
- Induced dipole + induced dipole (weakest)

Bonding Forces vs. Nonbonding (Intermolecular) Forces
- Bonding: Stronger attraction; generally higher energy compared to nonbonding.
  - Ionic & covalent
- Nonbonding are weaker and vary in strength based on molecular structure and types of interactions.
  - Ion-dipole
  - H bond
  - dipole-dipole
  - dipole-induced dipole
  - dispersion (london)

Electronegativity is the ability of an atom to attract electrons in a bond.
- increases with decreasing size
Bond Polarity: Differences in electronegativity lead to dipoles; unequal sharing of electrons creates polar covalent bonds.

Molecular shape determines overall polarity; polar bonds don’t always result in polar molecules if dipoles cancel (e.g., CO2).
A molecule is polar if it has at least one polar bond and does not have its dipoles cancel.

Solutes dissolve in solvents when their intermolecular forces are similar.
- “Like dissolves like” principle.
- Solvation (or hydration when water is the solvent) occurs when solvent molecules surround solute particles.

ΔH_sol’n = ΔH_solute + ΔH_solvent + ΔH_mix
- Positive indicates an endothermic process (energy absorbed);
- Negative indicates an exothermic process (energy released).

Solvation is the process of surrounding a gaseous solute particle with solvent particles.
When water is the solvent, solvation is called hydration.
ΔH_hydration = ΔH_solvent + ΔH_mix
Increase in charge = increase ΔH_hydration
Decrease in radius = increase ΔH_hydration
- Increase in charge & Decrease in radius = charge density
- charge/volume or charge/area

Vapor pressure increases with temperature; more molecules escape the liquid phase at higher temperatures.
- ↓ T ⇒↑ P_vap
Weaker intermolecular forces lead to higher vapor pressures.
- ↓ attraction ⇒↑ P_vap

Phase diagrams depict states and transitions (e.g., solid, liquid, gas) under different conditions of pressure and temperature.
- Triple point is where all three phases coexist in equilibrium.
- Critical point marks where the densities of liquid and gas become equal.

Two types: crystalline (orderly structure) and amorphous (disordered structure).
Unit cells are the basic building blocks of crystalline solids.
- Cubic lattice systems include types like simple cubic, body-centered cubic (bcc), and face-centered cubic (fcc).

Individual atoms, ions, or molecules are held together by different types of forces depending on the solid's composition.