CHEMISTRY 120
Fall 2009


Classes meet Monday, Wednesday, Thursday, Friday 11:10-12:20 Room 104 SC

Labs meet:

Nancy H. Kolodny

Paul Reisberg (laboratories)

Amanda McCarthy (laboratory)

FirstClass Conference: CHEM120-01-F09

Textbook: Steven S. Zumdahl, Chemical Principles, 5th Edition

Chemistry 120 is a one semester course for students with especially strong backgrounds in chemistry. It replaces Chemistry 105 and 205 as a prerequisite for upper level chemistry courses. Chemistry 120 includes both microscopic and macroscopic aspects of chemistry. The microscopic topics include the atomic nucleus, the behavior of electrons in atoms, and several approaches to molecular structure and bonding such as VSEPR, valence bond and molecular orbital models. Periodic properties of the elements are considered, with an emphasis on transition metals. Macroscopic topics include chemical kinetics and thermodynamics, electrochemistry, and chemical equilibrium as applied to acids, bases and complexes. The laboratory portion of the course complements and reinforces lecture material, and introduces statistical analysis of data, molecular modeling and computational chemistry.

Lectures and discussions will take place during the four periods assigned to the course each week. During some weeks class will meet three periods instead of four, as indicated on the attached schedule. Review sessions will be held during class periods at the end of each major unit of the course, or as needed. Classes will ordinarily be 70 minutes in length, but may be shorter.

After each class (before 6 AM the day of the next class) you are expected to submit a “one-minute paper” to our FirstClass Conference. The one-minute paper’s purposes are to help you identify the central theme(s) of the class and to point out to me what you found most confusing. I will then follow up on points of confusion, either individually (through e-mail or a meeting), through our FirstClass conference, or in the next class. You may contact me via e-mail or the FirstClass conference at any time with questions about lecture material. My posted office hours may be supplemented by appointments at other times during the week.

Problem sets will be assigned regularly. You are strongly encouraged to work with other students on them. Several problems from each set will be graded. In-class quizzes, midterm exams and the final exam will be based on material introduced in lecture and laboratory, and reviewed in problem sets. Individual requests for rescheduling of quizzes and midterm exams will be considered only if they are made prior to the quiz or exam.

The grade in Chemistry 120 will depend on all these components as follows:

In-class exams 28%
Final exam 20%
Quizzes 8%
Problem sets 10%
One minute papers 4%
Laboratory 30%

Note: Satisfactory completion of the laboratory portion of the course is necessary.
Any missed labs must be made up.

Your final grade will reflect both the numerical results of the above and an assessment of your progress throughout the semester.

The three in-class exams will be on Friday, October 9, Friday, November 6 and Friday, December 4. Half-hour quizzes will be announced two classes prior to when they will be given. A review session will be held in the class preceding each exam or quiz.

TENTATIVE SCHEDULE

Lecture Date Topic Zumdahl Chapter
1 9/9 Introduction to course; Nuclear Chemistry 21
2 9/10    
3 9/11    
4 9/14 Quantum Mechanics and Atomic Theory 2 (review), 12
5 9/16    
6 9/17    
7 9/18    
8 9/21    
9 9/23    
10 9/24    
11 9/25    
  9/28 No class  
12 9/30    
13 10/1 Bonding 13 (review), 14
14 10/2    
15 10/5    
16 10/7    
17 10/8 Review session  
18 10/9 Exam 1  
19 10/14 Monday schedule - class meets  
20 10/15    
21 10/16 Transition Metals and Coordination Chemistry  
22 10/19    
23 10/21    
24 10/17 Chemical Kinetics 15
25 10/23    
26 10/26    
27 10/28 Chemical Thermodynamics 9,10
28 10/29    
29 10/30    
30 11/2    
31 11/4    
32 11/5 Review Session  
33 11/6 Exam 2  
34 11/9    
35 11/11 Electrochemistry 11
36 11/12    
37 11/13    
38 11/16    
39 11/18 Properties of Solutions 17
40 11/19    
41 11/20 Chemical Equilibrium;Acids, Bases and Salts 6 (review), 7, 8
42 11/23    
43 11/25    
  11/26 Thanksgiving break - no class  
  11/27 Thanksgiving break - no class  
44 11/30    
45 12/2    
46 12/3 Review session  
47 12/4 Exam 3  
48 12/7    
49 12/9    
50 12/10    
51 12/11 Last class  

CHEMISTRY 120 OUTLINE
Fall 2009

  1. Nuclear Chemistry
    1. Historical Background
    2. Nature of the Nucleus
      1. Composition: nucleons
      2. Size
      3. Shape
      4. Mass
      5. Structure
      6. Energetics
    3. Nuclear Reactions
      1. Spontaneous decay
      2. Definition of radioactivity
      3. Particles/em radiation
      4. Units of radioactivity
      5. Examples and rules
      6. Kinetics of nuclear decay (see pp. 715-722)
    4. Applications of Nuclear Reactions
      1. Nuclear transformation by bombardment
      2. Nuclear fission: weapons and power plants
      3. Nuclear fusion

  2. Quantum Mechanics and Atomic Theory
    1. The Periodic Table
      1. History
      2. Atomic weights (masses)
    2. The Nuclear Atom: Building the Model
      1. 1885 Atomic Spectra: Rydberg Equation
      2. 1904 Thomson "Plum Pudding Model"
      3. 1911 Rutherford Nuclear Model
    3. The Nuclear Atom: The Quantum Model
      1. 1900 Planck's Quantum Hypothesis
      2. 1905 Einstein and Photons
      3. 1913 The Bohr Model
      4. 1924 de Broglie Hypothesis: Matter Waves
      5. 1925 Heisenberg Uncertainty Principle
      6. 1925-26 Schrödinger Quantum Mechanics: the hydrogen atom
        1. Quantization of energy
        2. Quantization of spatial distribution
        3. Electron spin quantum number
        4. Pauli Principle
        5. Examples
        6. Physical interpretation of Ψ
      7. Polyelectronic atoms
        1. Aufbau principle: building up atoms
        2. Hund's rule
        3. Orbital energies
          1. Energy and quantum number n
          2. Energy and quantum number l
            1. Effective nuclear charge Zeff
            2. Penetration
            3. Shielding
            4. Valence electrons
      8. Periodic Properties
        1. Atomic and ionic size (radius)
        2. Ionization Energy
        3. Electron Affinity

  3. Bonding
    1. Evidence for the existence of molecules
    2. Properties of molecules
      1. Chemical composition
      2. Bond lengths
      3. Bond energies, or strengths
      4. Molecular geometries
      5. Polarity (dipole moments)
      6. Magnetic properties
    3. Types of Bonds
      1. Ionic
      2. Covalent
      3. Metallic
      4. Intermolecular forces
    4. Valence Shell Electron Pair (VSEPR) Model for Molecular Geometries
    5. Localized Electron (Valence Bond) Model
      1. Bond formation: lengths, strengths
      2. sigma bonds
      3. pi bonds
      4. Molecular geometries
        1. hybridization
          1. rules for constructing hybrids
          2. types of hybrids
            1. sp
            2. sp2
            3. sp3
          3. multiple bonds
          4. what's missing: delocalization
    6. Molecular Orbital Model
      1. Basic assumptions
      2. Applications
        1. Homonuclear diatomic molecules
        2. Heteronuclear diatomic molecules
        3. Polyatomic molecules

  4. Transition Metals and Coordination Chemistry
    1. Historical Background
    2. Transition Metal Complexes
      1. Definitions and Nomenclature
        1. transition metal complex
        2. ligand
        3. coordination number
        4. isomers
          1. structural isomers
          2. stereoisomers
      2. Theory of Bonding for Transition Metal Complexes
        1. Requirements of a good theory
        2. Bonding theory: Crystal field model
      3. Biologically Significant compounds

  5. Chemical Kinetics
    1. History
    2. Definitions
    3. Rate Laws for Simple Reactions
      1. Zero Order
      2. First Order
      3. Second Order
    4. Experimental Determination of Rate Laws
      1. Graphical Methods
      2. Method of initial rates
    5. Reaction Mechanisms
    6. Activation Energy and Temperature Dependence of Reaction Rates
    7. Catalysis

  6. Chemical Thermodynamics
    1. Definition of Important Terms
    2. Heat and Work
    3. First Law of Thermodynamics ΔE = q + w
    4. Calorimetry
    5. Enthalpy H = E + PV
    6. Hess's Law
    7. Bond Energies
    8. Entropy and the Second Law
    9. Gibbs Free Energy ΔG = ΔH -TΔS
    10. Third Law of Thermodynamics
    11. Thermodynamics and Equilibrium

  7. Electrochemistry
    1. Redox Reactions
    2. Concepts of Electricity
    3. Electrochemical Cells
      1. Galvanic Cells
        1. Electrical Work
        2. The Standard Hydrogen Electrode
        3. Measuring equilibrium constants
        4. Concentration Cells
      2. Electrolytic Cells
        1. Writing the equations
        2. Calculating extent of reaction
      3. The Standard Hydrogen Electrode
    4. Practical Examples
      1. Lead Storage Battery
      2. Dry Cells
      3. Specialized Electrodes

  8. Properties of Solutions
    1. Formation of Solutions
    2. Colligative Properties
      1. Boiling Point Elevation and Freezing Point Depression
      2. Osmotic Pressure

  9. Chemical Equilibrium: Acids, Bases and Salts
    1. Historical Introduction and Definitions
      1. Etymology
      2. Phenomenological definitions
      3. Theoretical definitions
    2. Types of acids and bases; strengths
      1. Quantitative expression of acidity or basicity of solution: "p" notation
      2. Factors causing acid or base strength
        1. Binary acids
        2. Oxo-acids
        3. Bases
        4. Summary of criteria for acid and base strength
    3. Equilibrium in Acid-Base Solutions and pH calculations
      1. Strong Acids and Bases
      2. Weak Acids and Bases
        1. Weak monoprotic acids
        2. Weak bases
        3. Weak polyprotic acids
      3. Salts that act as acids or bases
      4. Buffers
        1. Effects of common ions
        2. Examples of Biological Buffers
        3. Calculation of pH of a buffer solution:
          Henderson-Hasselbalch equation
        4. Buffer Capacity
    4. Titrations and pH curves
      1. Definitions
        1. Titration
        2. Equivalence point
        3. Endpoint
      2. Calculation of a titration curve: weak base with a strong acid
        1. No acid added: we have a solution of a weak base.
        2. Not enough acid added to neutralize base: buffer region
        3. Exactly enough acid added to neutralize base: equivalence point
        4. Excess acid added beyond equivalence point
      3. Indicators
    5. Solubility Equilibria
      1. Ksp
      2. Common Ion Effect