Chemistry 111             Spring 2005     Review for the Final Exam     Dr. Mahler


The Final Exam is a standardized multiple choice exam from the American Chemical Society. More information on the exam is available in “Preparing for Your ACS Examination in General Chemistry: The Official Guide”, on two-hour reserve in the Snowden Library.


Areas covered in ‘The Official Guide’ are as follows: Atomic Structure, Molecular Structure and Bonding, Stoichiometry, States of Matter / Solutions, Energetics, Dynamics, Equilibrium, Electrochemistry / Redox, Descriptive Chemistry / Periodicity, and Laboratory Chemistry.


What follows is a chapter by chapter brief overview of the textbook, Hill and Petrucci, General Chemistry: An Integrated Approach, 3rd Ed.


Chapter One: Chmeistry: Matter and Measurement; Scientific Method; Composition of Matter (elements vs. compounds and homogenous vs. heterogeneous mixtures); Physical and chemical changes; SI (Metric system) units, prefixes, and conversions; Precision, accuracy and significant figures; Calculations showing units; Density.


Chapter Two: Atoms, Molecules and Ions; Atomic Theory; Laws of conservation of mass, constant composition, and multiple proportions; Atomic particles especially protons, neutrons and electrons; Mass number A and atomic number Z; Isotopes; Chemical formulas: empirical, molecular, structural, and condensed structural formulas; Periodic Table; Names of ionic, molecular and simple organic compounds; simple acids and bases.


Chapter Three: Stoichiometry: Chemical Calculations; Molecular and formula masses; Moles, Avogadro’s number, molar / atomic mass; Mass percent; Empirical vs. molecular formulas; Balanced chemical equations; Stoichiometric calculations: limiting reactant, theoretical yield, actual yield, percent yield; Molarity, dilution, and solution concentrations.


Chapter Four: Chemical Reactions in Aqueous Solutions; Ionic compounds in solution: weak and strong electrolytes, non-electrolytes; Strong and weak acids and bases; Neutralization (acid/base) reactions and titrations; Solubility rules and precipitation reactions; Oxidation numbers and oxidation-reduction reactions; Oxidizing and reducing agents; Activity series of metals.


Chapter Five: Gases; Gas theory; Pressure, barometers and manometers; Pressure units (Pascal, atmosphere, torr or mm Hg); Other gas properties: volume, moles, (absolute) temperature; STP and molar volume; Various gas laws and the Ideal Gas Law, PV = nRT; Stoichiometric, Molecular mass, and density calculations for gases; Dalton’s Law of Partial Pressures; Kinetic-molecular theory; Real gases and van der Waals equation.


Chapter Six: Thermochemistry; Energy changes in physical processes or chemical reactions; System and surroundings; Kinetic, potential and internal energies; heat and work; Internal energy, U, and Enthalpy, H; State functions; Constant volume and pressure conditions; First Law of Thermodynamics; Thermochemical equations; endo- and exo-thermic; Types of calorimeters; Standard enthalpy change, ∆H°, and standard enthalpy of formation; Hess’ Law and Calculations – sum of products times their stoich. numbers minus the sum of reactants times their stoich. numbers.


Chapter Seven: Atomic Structure; Electrons, cathode rays, fundamental charge; Thomson and Rutherford’s atomic models; Nucleus (protons and neutrons) surrounded by electrons; Electromagnetic radiation (light, etc.), frequency, wavelength, and speed; Visible light and its properties; Quantum – energy is limited to certain values; Bohr’s theory of hydrogen atom; Wavelike nature of small particles; Uncertainty principle; Wave equation as atomic model for electrons; Quantum numbers n, l, ml, and ms; Atomic orbitals: s and p.


Chapter Eight: Electron Configurations, Atomic Properties, and the Periodic Table; Wave mechanics of the hydrogen atom (one electron) can be extended to multielectron atoms; differences – energy levels are lower and split (compared to H atom); Shells (same n) and subshells (s, p, d, f orbitals); Electron configuration: spdf notation and orbital diagram; Filling electrons in orbitals: lowest energy filled first, Pauli exclusion principle, Hund’s Rule, and Aufbau principle; Periodic table structure: blocks (spdf), periods (rows), and groups (columns); Diamagnetic vs. paramagnetic; Periodic trends in atomic properties: atomic radius, ionic radius, ionization energy, and electron affinity; Metals, non-metals, metalloids, and noble gases.


Chapter Nine: Chemical Bonds; Lewis symbols; Ions: Lattice energy and the Born-Haber cycle; Covalent bonds, Lewis structures, octet rule; Electronegativity and polarity of bonds; Types of bonds: coordinate covalent, single, multiple; How to write Lewis dot structures; Resonance and hybrids; Exceptions to the octet rule; Bond order (single, double, triple) and bond length, bond dissociation energy; simple organic reactions involving multiple bonds.


Chapter Ten: Bonding Theory and Molecular Structure; Molecular / polyatomic ion geometry and shape from VSEPR theory; bonds versus lone pairs; polar covalent bond theory, bond dipoles and molecular polarity (polar or non-polar); Valence bond theory, overlap of atomic orbitals; hybridization, and hybridized orbitals (agreement with VSEPR);  multiple bonds and unhybridized orbitals; sigma and pi bonds; Molecular orbital theory; atomic orbitals combine to form molecular orbitals; bonding and anti-bonding orbitals and bond order; combinations of VB and MO theory.


Chapter Eleven: States of Matter and Intermolecular Forces; States of matter; Phase changes; vaporization / condensation; vapor pressure, normal boiling point, critical point and temperature; Phase diagrams: solid, liquid, gas and vaporization, sublimation and fusion curves; Triple point; Intermolecular (van der Waals) forces: dispersion forces, dipole-dipole forces, and dipole-induced dipole forces; Hydrogen bonds, nature and limitations; Viscosity and Surface tension (End of Chemistry 110);


Chapter Eleven:  (Start of Chemistry 111) Solids: covalent network vs. molecular solids, ionic solids; Crystal structures and three types of cubic lattices; unit cells and counting atoms / ions in them; Closest packing of spheres, hcp, ccp, holes in structures.


Chapter Twelve: Physical Properties of Solutions; Solutions and how to describe them: molarity, molality, mole fraction (also percentage), parts per million (ppm), …billion (ppb), and …trillion (ppt); ideal and no-ideal solutions and intermolecular forces (homogeneous vs. heterogeneous mixtures); Solubility and saturated solution, its temperature dependence; Raoult’s and Henry’s Laws; fractional distillation; Colligative properties: freezing point depression, boiling point elevation, and osmotic pressure; van’t Hoff factor, i, for colligative properties; colloids and their properties.


Chapter Thirteen: Chemical Kinetics: Rates and Mechanisms of Chemical Reactions; Rates of reactions, units; Rate law, rate constant, order in individual species and overall; Integrated rate laws: first, second and zeroth order, half lives, and units of k; Molecular collisions, transition state,activation energy and enthalpy; Temperature dependence of k and the Arrhenius equation; Reaction mechanisms and rate laws; Catalysts and enzymes.


Chapter Fourteen: Chemical equilibrium; equilibrium expression; K in terms of concentration and pressure and interconversions; size of K and completeness of reaction; reaction quotient Q, predicting direction of reaction in non-equilibrium conditions; Le Chatelier’s principle: response of the system to changes in amounts of reactants or products, volume and pressure, or temperature; catalysts; types of equilibrium calculations.


Chapter Fifteen: Acids, Bases and Acid-Base equilibria; Arrhenius and Bronsted-Lowry acid-base theories; conjugate acid-base pairs and their equilibrium reactions; Relative strength of acids and bases and their conjugates; Water, amphiprotic, self-ionization, Kw, pH and pOH; pH scale; Hydrolysis reactions of acid-base salts; Common Ion effect; Buffers and their equations; Indicators and titrations; Lewis acid-base theory.


Chapter Sixteen: More Equilibria: Slightly Soluble Salts and Complex Ions; Solubility product, Ksp, and molar solubility; Common ion effect and completeness of precipitation; Predicting whether precipitation will occur; pH dependence of some solubilities; Complex ions, ligands, Lewis acid-base reactions, Kf; Increased solubilities by complex formation; Acidic character of some complex ions and amphoteric oxides and hydroxides.


Chapter Seventeen: Thermodynamics: Spontaneity, Entropy and Free Energy; Spontaneity and criteria for it; Entropy, definition, total, system and surroundings; Second and Third Laws of Thermodynamics; Gibbs Free Energy change, ∆G, and Standard Gibbs Free Energy change, ∆G°, and standard Gibbs Free Energy change of formation; Calculations – sum of products times their stoich. numbers minus the sum of reactants times their stoich. numbers; Relations to Keq; van’t Hoff equation – variation of K with temperature; coupled reactions.



Chapter Eighteen: Electrochemistry; Oxidation and reduction reactions and half reactions; Electrochemical cells, half cells, anode, cathode, salt bridge; Cell diagrams; Standard potentials (written as reductions) and how to calculate standard cell potentials; Relating Potential (E) to Gibbs Free Energy and Keq; Nernst equation – how changes in concentration affect cell potentials; pH meters. (through 18.6, not sure how much further we’ll get, we’ll at least talk about batteries and electrolytic cells).


Chapter Nineteen: I hope to at least briefly go over some basics of nuclear chemistry.


Equations used – I quickly looked through the practice questions in ‘The Official Guide’ and here are the equations / formulas used in each section which would likely have been given to you on a quiz or exam:


Atomic Structure (weighted average for atomic mass from isotopes),

Molecular Structure and Bonding (none but lots of VSEPR, Lewis structures, etc.),

Stoichiometry (none but lots of stoichimetry calcs, yields, limiting reactant),

States of Matter / Solutions (ideal gas law, molarity, dilution, mole fraction, f.p depression given the constant),

Energetics (q lost = q gained, q = mc∆T, density, ∆H (products minus reactants), general enthalpy calcs).

Dynamics (method of initial rates, first order rate law and half-life, knowing dependence of zero (conc. vs. time) and second order (1/conc. vs. time) rate laws), Note - they give the Arrhenius equation,

Equilibrium (none – they give values for Ksp, Ka, Kb),

Electrochemistry / Redox (E cell = E right or red minus E left or ox), Note – they give the Nernst equation,

Descriptive Chemistry / Periodicity (none),

Laboratory Chemistry (dilution).