Ch 10 and 11

Intramolecular forces

• Covalent bonds within a molecule

• Strong

• Strength determines chemical properties

Intermolecular forces (IMF’s)

• Weak attraction forces between molecules

• Strength determines Physical properties of the substance

Types of IMF’s

1. London Dispersion Forces (LDF’s) (weakest)

• Repulsion between electrons of atoms cause electron cloud to polarize

• Instantaneous dipoles form as a result

• Present in all molecules and ions

• Larger electron clouds, more atoms in a molecule, and increased surface area cause stronger LDF’s because the electron clouds are more polarizable because they are further from the nucleus or there are more places for LDF’s to form

  • Chain like molecules have stronger LDF’s than sphere-like molecules

2. Dipole-Dipole Attractions (3rd strongest)

• Occurs between polar molecules

• Attractions > Repulsion

• Hydrogen bonding (2nd Strongest)

  • Special type of D-D interaction

  • FON bonds are very polar, very strong D-D forces

3. Ion-Dipole Attractions (Strongest)

• Attractions between ion and charged end of polar molecules

• In water, negative oxygen side of H₂O would be attracted to cation, and positive H to the anion

• Ion-Induced Dipole Attractions

  • Attractions between ion and dipole it induces on neighboring molecules

  • Depends on ion charge and polarizability on neighbor

Physical Properties that depend on how tightly molecules pack

• Compressibility

  • Gases → Highly compressible

  • Solids and liquids → Nearly incompressible

• Diffusion

  • Rapid in gases

  • Slower in liquids

  • Nearly none in solids (increases with temp increase)

• Phase

  • Solids → Strongest IMF, Gases → Weakest IMF

Surface Tension

• Measure of the amount of energy needed to expand the surface area of a liquid

• Less IMF’s at surface of liquid, PE is greater

• Molecules form spheres to minimize surface area and therefore PE (water droplets on leaves)

Wetting

• The spreading of a liquid across a surface to form a thin film

• IMF’s between surface and liquid have to be as strong as IMF’s in liquid

• Surfactants lower surface tension, making water wetter, and allows polar water to spread across nonpolar greasy glass by being polar and nonpolar

Viscosity

• A liquids resistance to flow

• As size increases and temp decreases, viscosity increases

Vapor Pressure

• Pressure molecules exert when they evaporate

• Pressure of gas when they are at equilibrium with liquid or solid phase

Le Chatelier’s Principle

• When an equilibrium is upset, it will try to restore equilibrium

Types of Solids

Crystalline Solids → Repeating structure patterns i.e. salt, sugar, quartz, metals

• Ionic Crystals

  • Cations and Anions

  • Hard

  • High melting points

  • Only conduct electricity when liquid

  • Brittle

• Molecular Crystal

  • Lattice sites are occupied by atoms or molecules

  • Soft, low melting point, do not conduct

  • Vander walls forces

• Covalent Crystals/Network Solids

  • Lattice positions occupied by atoms that are covalently bonded to other atoms at neighboring lattice sites

  • Hard, high melting points, poor conductors

• Metallic crystals

  • Cations surrounded by sea of electrons

  • Conduct heat and electricity

  • Luster

• Metal Alloys → Mixture of metals

  • Substitutional Alloy → some metals are replaces by other metals of similar size

  • Interstitial Alloy → Some of the holes (interstices) of metal structures are occupied by small atoms

Amorphous Solids → Do not have repeating Patterns

Solutions

• Homogenous mixtures of solvent and solute

• Gases mix spontaneously

• Liquids depend on IMF’s and entropy to mix

  • More similar solute and solvent IMF’s, the more they mix, “like dissolves like”

• Miscible → 2 solutions that are soluble in each other

• Immiscible → Do not mix, form 2 phases

• Solids in Liquids

  • NaCl dissolves in water because ion-dipole forces of water with ions are strong enough to overcome ion-ion attractions

  • Ions become hydrated when surrounded by water molecules

Heat of Solution

• Total energy absorbed or released when a solute dissolves in a solvent at constant pressure

• Lattice energy → Energy released when ionic crystals form

• Hsol = Hlatt + Hhydr

Henry Law

• Cgas =KhPgas

• C1/P1 = C2/P2

Vapor Pressure

• nonvolatile solute lowers the vapor pressure of a solvent because the number of solvent molecules at the surface is lower because there are more solute molecules

IMF’s are stronger when the solute and solvent are separate, the reaction would be endothermic because energy would have to be added in order for them to form less efficient bonds and visa versa