Exam 1 Review

Lecture 1: Forces in Molecules

• Difference between Intramolecular and Intermolecular Forces:

  • Intramolecular Forces:

  • Forces that hold atoms together within a molecule.

  • Types include:

    • Covalent bonding

    • Ionic bonding

    • Metallic bonding

  • Generally stronger than intermolecular forces (IMF).

  • Intermolecular Forces (IMF):

  • Attractions between molecules.

  • Weaker than intramolecular forces with varying strengths.

  • Strength hierarchy:

    • London Dispersion Forces < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole < Covalent < Ionic < Metallic

  • Stronger IMF lead to:

    • Higher melting/freezing/boiling points

    • Higher viscosity

    • Greater surface tension

    • Higher enthalpies of fusion, vaporization, sublimation

    • Lower vapor pressure

• Important Concepts:

  • Lewis Structure: Key to understanding molecular structure and behavior.

Types of Intermolecular Forces

• London Dispersion Forces:

  • Weakest IMF, present in all molecules.

  • Caused by temporary dipoles due to electron movement.

  • Strength increases with molecular weight due to greater polarization potential.

• Dipole-Dipole Interactions:

  • Occur between polar molecules.

  • Positive end of one molecule attracts the negative end of another.

• Hydrogen Bonds:

  • Special case of dipole-dipole interaction.

  • Involves hydrogen atoms bonded to highly electronegative atoms (N, O, F).

  • Stronger than regular dipole-dipole interactions.

• Ion-Dipole Forces:

  • Interactions between an ion and a polar molecule, crucial in solvation.

• Collective Terms:

  • Dipole-induced dipoles and London Dispersion forces are referred to as Van der Waals forces.

  • Instantaneous dipoles can form when nonpolar molecules are close together.

Lecture 2: Heat and Phase Changes

• Phase Diagrams: Understand how temperature and pressure affect phases.

• Endothermic vs Exothermic Processes:

  • Endothermic: Absorbs heat. e.g., Solid → Liquid → Gas (bond breaking).

  • Exothermic: Releases heat. e.g., Gas → Liquid → Solid (bond forming).

• Clausius-Clapeyron Equation:

  • R = 8.314 ext{ J/mol*K}

  • Pressure (P) is directly related to temperature during phase changes.

Cooling & Heating Curves

• Temperature Change:

  • Q=mc\Delta T

  • where:

    • m = mass

    • c = specific heat

    • \Delta T = final temp - initial temp

• Phase Change:

  • Q=m\cdot\Delta H

  • Remember the signs:

  • Endothermic: positive \Delta H

  • Exothermic: negative \Delta H

  • Types of \Delta H : Fusion, Evaporation, Sublimation.

Lecture 3: Unit Cells

• Volume Calculation:

  • V=\left(edge\cdot leng\operatorname{th}\right)^3\left(a^3\right)

  • 1 Å = 10^{-10} m

  • 1 m = 10^{12} pm

Lecture 4: Solution Process

• "Like Dissolves Like" Principle:

  • Polar solvents dissolve polar solutes, nonpolar solvents dissolve nonpolar solutes.

  • Requires breaking solute-solute and solvent-solvent interactions.

• Enthalpy of Solution:

  • riangle H{soln} = riangle H1 + riangle H2 + riangle H3

Henry's Law

• Solubility Factors:

  • Solids: Higher temperature increases solubility.

  • Gases: Higher pressure increases solubility; lower temperature helps gas solubility as well.

• Henry’s Law Equation:

  • C=kP

  • where:

    • C = concentration of gas (M)

    • P = partial pressure (atm)

    • k = Henry’s law constant (M/atm).

Lecture 5: Concentration

• Molarity (M):

  • =molesofsolute/litersofsolution

• Molality (m):

  • =molesofsolute/kilogramsofsolvent

• Volume Percent (%v/v):

  • =volumeofsolute/volumeofsolution*100

• Weight Percent (%w/w):

  • =massofsolute/massofsolution*100

• Weight by Volume Percent (%w/v):

  • =massofsolute/volumeofsolution*100

• Mole Fraction (X):

  • =molesofsolute/molesoftotalsolution

• ppm and ppb:

  • ppm=milligramsofsolute/litersofsolution

  • ppb=microgramsofsolute/litersofsolution

Lecture 6: Colligative Properties

• Types:

  • Vapor pressure lowering

  • Boiling point elevation

  • Freezing point depression

  • Osmotic pressure

• Calculations for water:

  • New boiling point: 100+\Delta T

  • New freezing point: 0-\Delta T

• Van 't Hoff Factor (i):

  • Number of particles a solute breaks into upon dissolving.

• Effect of Concentration:

  • Higher concentration leads to higher boiling points, lower freezing points, lower vapor pressures, and higher osmotic pressure.

Lecture 7: Colloids

• Definition:

  • A suspension of tiny particles in a medium (not dissolving).

  • Intermediate between homogeneous and heterogeneous mixtures.

  • Particle size: 10^3 to 10^6

• Coagulation:

  • Process where a colloid is destroyed, typically by heating or adding an electrolyte.