TO STUDENTS:

        Most of the lecture will be from this HTML notes and several PowerPoint Lecture notes. Some of the material  in this HTML notes are based on "General Chemistry" (7th Ed., Ebbing & Garmon), the old textbook used  prior to 2004.

        Two PowerPoint Lecture Notes will complement this HTML notes: the first one is a modified version of the Power Point Slides (by Prof. J. D. Robertson) an old textbook “Chemistry( 8th Ed., Raymond Chang)” , the second one being the Power Point Slides (by Prof. J. D. BookStaver) that came with the current textbook “Chemistry - The Central Science" (10th Ed., Brown, LeMay & Bursten).

        50%  of the lecture material will be selected from this HTML notes, another 50% will come from the Power Point Slides.   In addition, these will be reinforced  with some additional board work.

Textbook used:

      General Chemistry (7th Ed., Ebbing & Garmon): 1992 ~ Summer 2004
      Chemistry (8th Ed., Raymond Chang), Fall 2004 ~ Fall 2006

       Chemistry (10th Ed., Brown, Lemay & Bursten), Fall 2006 ~

PowerPoint Lecture Slides (BLB):   Chap. 1

PowerPoint Lecture Slides (BLB):   Chap. 2

PowerPoint Lecture Slides (BLB):   Chap. 3

PowerPoint Lecture Slides (BLB):   Chap. 4

PowerPoint Lecture Slides (BLB):   Chap. 5

PowerPoint Lecture Slides (BLB):   Chap. 6    

PowerPoint Lecture Slides (BLB):   Chap. 7

PowerPoint Lecture Slides (BLB):   Chap. 8  

PowerPoint Lecture Slides (BLB):   Chap. 9

PowerPoint Lecture Slides (BLB):   Chap. 10

CHAPTER 1 , 2 , 3 , 4 , 5 , 6 , 7 , 8 , 9 , 10 

( * Additional subjects not given in the Text) 

Chapter 1  CHEMISTRY and MEASUREMENT: Introduction
CH 1  Chemistry and Measurement

1. The Science of Chemistry:
      Chemistry as the Central Science
      Basis for other Sciences, Engineering and Technologies
      Importance of Chemistry : Chemistry/Chemical Reactions are everywhere.
      Chemistry and Matter

2. The Scientific Method: Experiment and Explanation ( Chart )
     Observations, Hypothesis, Law, and Theory

3. Law of Conservation of Mass:
    What is a Natural Law?
     Some Important Laws in Chemistry

4. Matter : Physical State and Chemical Constitution:
     Three Forms of Matter: Solids, Liquids, and Gases (Phase Changes)
     The Effect Temperature on Matter *
     Chemical vs. Physical Changes: Chemical vs. Physical Properties
     Intensive vs. Extensive Properties

• Some Basic Definitions/Terms :

Methods of Chemical Separation from a Mixture:
     Examples>  Distillation: Petroleum Refining
           Chromatography : Polar vs. Nonpolar Compounds

5. Measurement, Significant Figures, and Errors/Uncertainty
   Your Weight
     The Rules, Scientific Notation
    Significant Figure in Addition/Subtraction & Multiplication/Division
        Measured Numbers vs. Exact Numbers: Examples

   Introductory Section for Basic Statistics (Optional Material):
      Accuracy vs Precision : Examples & Calculations (Printed Version)
  Calculation of Average & Std. Deviation [More Exercise to come in the Lab]

6. SI Units:
    Base Units & SI Prefixes
    Ranges of Length, Mass & Time

7. Derived Units :
      Area, Volume, Density (d=m/V), Speed, Acceleration,
      Force, Pressure, Momentum, Energy/Work, Power [Board work]

8. Unit Conversions & Dimensional Analysis(Factor-Label Method)
    Conversion Examples : Auto Speed , Ocean Volume,
             Annual  Food/Drink Consumption,  More Conversion Exercise
      Temperature Conversion (w/ Clear Background ): Thermometer
           ^oF = (9/5) ^oC + 32
9. Propagation of Random Errors* - Moved to CHEM 1212L -
 

Chapter 2 ATOMS, MOLECULES, and IONS

1. Atomic Theory of Matter (alternate slide ):
     Democritus (~ BC 400), Dalton's Postulate (1805),
     Law of Multiple Proportion , Atomic Symbols (See the Table)

2. Structure of the Atom:
      Discovery of the Electron: Thomson's Cathode Ray Expt. (1897; NP, 1906)
                     e/m  = 1.76x10⁸ C/g
      Electronic Charge : Millikan's Oil Drop Experiment   ( Figure ) (1909; NP, 1923)
             e = 1.6021x10^-19 C,      m_e = 9.109 x 10^-31 kg
      Model of the Atom : Thomson's Plum Pudding Model : defunct
      Rutherford's Gold Foil Experiment (1906-1911; NP, 1908)
           - Most of the particle would be deflected if nuclei were large.
3. Nuclear Structure; Isotopes
      Electron, Proton (Rurherford, 1911), Neutron(Chadwick, 1932; NP, 1935)
      Atomic Number (Z) , Mass Number (A):  A = #p + #n,  Z = #p = #e
           & Nuclide Symbol
      Subatomic Particles: Summary Table

4. Atomic Weights: Atomic Weight of an Elements
      Relative Atomic Masses, Mass Spectrometry , Example with Neon
           1amu = 1.6605x10^-27kg
5. Periodic Table of the Elements:Overview(4-1, Whitten et al)
     Mendeleev & Lothar Meyer, Period & Groups,
     Metals, Nonmetals, Metalloids: General Properties
     Main Group (A Group) vs. Transition Elements (B Group)
     Metal Plus non-Metal

6. Chemical Formula ( Empirical Formula vs. Molecular Formula )
     Types of Substances : Molecular vs. Ionic Substances
                               (Covalent vs. Ionic Bonds)
   Formation of NaCl: Movie
     Molecule, Molecular Formula, Structural Formula & Molecular Model
       [Refer to the Figure ]
     Ions & Ionic Compounds : Cations vs. Anions

7. Organic Compounds : "Molecular substances that contain Carbon"

8. Naming Simple Compounds:
   (A) Ionic Compounds : Naming Ions, Polyatomic Ions
      Common Cations & Anions (Tables)
            Rules for Charge on Monoatomic ions
   (B) Molecular Compounds: Binary Compounds, Greek Prefixes
      Acids & Corresponding Ions (Oxoacids & Oxoanions )
           Hydrates, Greek Prefixes for naming Compounds (Table)

       Review & Other Examples

9. Writing Chemical Equations :
       represent a chemical reaction in terms of chemical formulae of reactants & products.

10. Balancing Chemical Equations: Why needed?
     (A) Simple Inspection Method ( altenate slide );
               Example: Combustion of Butane & Gasoline
   *(B) Algebraic Method : Example # 2
           Matrix Method (for advanced students)
 

- FIRST CLASS EXAM - Covers Chap. 1 & 2: Sample Test #1

Chapter 3 Calculations with Chemical Formulas and Equation
- Stoichiometry

0. Introduction : ( alternate slide )

1. Molecular Weight ( alt slide ) and Formula Weight

2. The Mole Concept ( alt slide ):  Big Numbers
      Avogadro's Number:  6.022 x10²³ = 1 mole
      Molar Mass & Mole Calculations:   # moles  = mass/molar mass
                                                                    i.e.,         n = m/M_m
      Examples: ZnI ₂

3. Mass Percentages ( Composition in Mass %) from the Formula
      Examples: #1 Formaldehyde, #2 Ferric Sulfate

4. Elemental Analysis; Experiment
      Percentages of Carbon, Hydrogen & Oxygen
      Examples: Acetic Acid

5. Determining Formulas
      Empirical Formula & Molecular Formula
      How to find Empirical Formula/Molecular Formula from mass data ?
      Examples: Acetic Acid , Citric Acid in Lemon, Dry Cleaning Solvent

6. Molar Interpretation of Chemical Reaction
     Examples: Synthesis of Ammonia , Combustion of Butane

7. Amounts of Substances in a Chemical Reaction
      Examples: Hematite , Calcium Hypochlorite (Bleaching Agent)
           How much O₂ is needed to burn off 1 lb of Fat ?
       How much O₂ is needed to burn off 1 lb of Sugar (a Carbohydrate) ?
8. Limiting Reagent: Analogy
     Examples: ZnCl ₂ , Acetic Acid , Sulfuric Acid
     Theoretical and Percentage Yields: Roasting Zn Ore

9.  Other Stoichiometric Problems (Optional)
    Volume Relationship in Gaseous Reaction , Aroma of Wine

Chapter 4 CHEMICAL REACTIONS

1. Ionic Theory of Solutions: Arrhenius (1884; NP, 1903)
    Electrolytes vs. Nonelectrolytes ,Common Electrolytes , (in Body )

2. Molecular and Ionic Equations
   Molecular Equations, Ionic Equations & Net Ionic Equations

TYPES of CHEMICAL REACTIONS (More Examples:1 ,2 ): Movie

   (1)  Combination Reactions:                     A + B       - >   AB
   (2)  Decomposition Reactions:                AB           - >   A + B
   (3)  Displacement Reactions:                   AB + C    - >  AC + B
   (4)  Metathesis (Exchange) Reactions:  AC + BD - >  AD + BC

3. Precipitation Reactions
     Soluble vs. Insoluble Salts: Solubility Rules

4. Acid-Base Reactions
      Acids & Bases , Acid-Base Indicators
      Strong Acids and Weak Acids, Strong Bases & Weak Bases
      Neutralization Reactions

5. Oxidation-Reduction Reactions: Redox Concept
    Examples: MgO , CaO
     Displacement Reactions: Dissolution of Cu , Activity Series
       Oxidation-Number Rules; Long, (Brief , Table)
       Oxidation Number (or Oxidation State): Main Group & Transition Elements
      Exercises : Why Redo Reactions Occurs? ( #1 , #2 , #3 )
       Combustion Reactions: " What is a Fire?"

6. Balancing Oxidation Reduction Reactions
         (1) Simple Inspection Method
       (2) Oxidation Number Method
       (3) Half Reaction Method
       (4) Algebraic Method : Common Method, Matrix Method

7. Solutions and Molar Concentration (Morality)
    Concentrations , % Concentrations Examples ,
      Converting Mass % to Mole % ,
      n, M & V Relationship : M=n/V, n=MV
      Eg., Sodium Nitrate , (alt slide)

8. Diluting Solutions:
    How to Prepare a Solution? ,   Dilution

9. Gravimetric Analysis:
    Example #1,
                  #2 Barium Chromate (Ba²⁺+CrO₄^2-_(yel) -> BaCrO_{4(yel ppt)})

10. Volumetric Analysis: Principle
     Examples of Acid-Base Titration:
      Acetic Acid with a Base (Commercial Vinegar ~ 5%, FDA)
           -> Board work for % Calculation
      H₂SO₄, HCl

11. Spectrophotometric Analysis*
        (to be covered in the Lab Class)

- SECOND CLASS EXAM - Covers Chap. 3 & 4: Sample Test #2

Chapter 5 THE GASEOUS STATE
 

States of Matter

1. Gas Pressure and Its Measurement
   Units of Pressure, mmHg(torr), atm , Pa
         1 atm = 760 mmHg
                 = 101,325 Pa
                 = 14.7 psi
   Pressure at a base:     p = gdh
         Example: Penny on a table, Pressure Conversion (Atm to inHg )
   Barometer, Manometer :
   Torricelli's Experiment , (alt: Atmospheric Pressure )

2. Empirical Gas Laws :
    Boyle's Law(Table , Curve ):  PV = Constant
    Charle's Law:                   V/T= constant
    Combined Gas Law :          PV/T = Constant
    Avogadro's Laws:              V = kn
      Exercise
   Mass of O₂
   Sip a Drink , Lung

3. The Ideal Gas Law : PV=nRT
   Gas Density: PV=(m/M_m)RT
                       m/V=PM_m/RT = density
      O ₂ ,
   Molecular-Weight Determination: M_m = mRT/PV
   Applications: Oxygen, Balloon

4. Stoichiometry problems involving with Gas Volumes
   Example: Air Bag

5. Dalton's Law of Partial Pressures ( alt. ) for Gas Mixture
      Total Pressure = Sum of all Partial Pressures
   Partial Pressure & Mole Fraction (X)
   Applications: Mole Fraction of N₂ in Dry Air
   Collecting Gases Over Water : Vapor Pressure of Water
   Victor-Meyer's Method for Determining Molecular Wt.
    Composition of Air : Inhaled vs. Exhaled, Daily O₂ Consumption
   Average Molar Mass of Air

7. Molecular Speeds: Graham's Law ( alt. ) Diffusion and Effusion
   Molecular Speed , Maxwell Distribution
   Root-Mean-Square (rms) Molecular Speed: v=(3RT/M)^1/2
   RMS Speed of N ₂ at room temperature, Mean-Free-Path Table
   Temperature from RMS

8. Real Gases vs. Ideal Gases ( alt .)
    van der Waals Equation (NP, 1910):
      (p+an²/v²)(P-nb)=nRT
    Application

9. The Atmosphere
   The Composition of Air

For Advanced Students : Pressure from Collisions , Universal Gas Constant

Chapter 6 THERMOCHEMISTRY

Overview
The Questions (Written , Heat Measurement )
Brief Summary *(Endothermic and Exothermic Processes)
Natural Change and Energy

1. Energy and its Units Work :   Energy Conversion Factor
   Types of Energy (alt: I , II ),
          Kinetic Energy and Potential Energies
          Internal Energy (U)
   Law of Conservation of Energy

2. Heat of Reaction
   System and Surrounding: Example , System Variables *
   Definition, Heat of Reaction, Heat Transfer (alt)

3. Enthalpy (H, Heat Content) and Enthalpy Change ( H)
   Enthalpy Diagram
   Pressure Volume Work;
   Enthalpy of Reaction, Enthalpy and Internal Energy
   Example (Sodium and Water: Figure )
   State Function: Figure

4. Thermochemical Equations
   Two Rules

5. Applying Stoichiometry to Heats of Reaction
   Example

6. Measuring Heats of Reaction: Measurement , Calorimeter
   Heat Capacity (C): q = C(t₂-t₁),  C = ms
   Specific Heat(s): q = ms(t₂-t₁)
   Example: Graphite
   Heat of Neutralization

7. Hess's Law
   Application , Enthalpy Diagram (A , B )
   Sublimation Energy Calculation (Ex 1, Ex 2)

8. Standard Enthalpy of Formation * (or Standard Heat of Formation)
   Standard State, The Table
   Standard Enthalpy of Reaction: Example ( Methane *)
   Answer to the Question

9. Fuels and Foods, Commercial Fuels, and Rocket Fuels (Chart )
      Energy Expenditure for Activities
   Foods as Fuels, Fossil Fuels, Coal Burning
   Coal Gasification and Liquefaction, Rocket Fuels
   Costs of Energy

House Heating Problem

Chapter 7 QUANTUM THEORY of the ATOM: An Intro

1. Electromagnetic Radiation: Wave Nature of Light
     Continuous Spectrum vs. Line Spectrum
     Emission Spectra of Hydrogen
    Emission Spectra: Finger Print of Atoms

   Waves : Sonic Wave , Electromagnetic Wave : Spectra
       Wavwlength (l), Frequency( n ), Velocity( c ): c = ln
        Calculation Examples
       3 Properties of Wave:  Reflection, Refraction, Diffraction

Old Physics (Galileo-Newton Mechanics) vs. New Physics
Three Paradoxes that couldn't be explained with the Classical Physics

2. Quantum Effects and  Photons

   Planck's Quantum Theory (1900; NP, 1918): E = nhn
   Einstein's Photon Theory (1905; NP, 1921): Photoelectric Effect: E = hn
          Energy of EMW, Planck's Constant: Example of Calculation

   Balmer's Formula (1885): the clue for the Puzzle of H Emission Spectra

3. The Bohr Model of the Hydrogen Atom (1913; NP, 1922): E = - R_H/n²
   Emission and Absorption of Light
   Transition Energy Calculation , Balmer Series etc.

4. Quantum Mechanics
    de Broglie's "Wavicle": Wave-Particle Duality (1924; NP, 1929): l = h/(mv)
      Electron Diffraction by Davisson/Germer (US) & G. Thomson (GB) (1927; NP, 1937)
          Electron Microscope by E. Ruska (1933, NP: 1986)
   Heisenberg's Uncertainty Principle (1927; NP, 1932)
   Schrodinger's Wave Equation (1926; NP, 1933)

5. Quantum Numbers and Atomic Orbitals
   (a) Principle Quantum Number (n):    n = 1,2,3,4 ^{. . .}
   (b) Angular Momentum Quantum Number (l): l = 0,^{. . .} ,(l-1)
   (c) Magnetic Quantum Number (m_l): m_l = -l,^{. . .},0,^{. . .},+l
   (d) Spin Quantum Number (ms):      m_s = +1/2, -1/2
   Table of Permissible Quantum Numbers ,: Exercise
   Shapes of Orbitals: s (2D , 3D ), p , d , f

Chapter 8 ELECTRONIC STRUCTURES of the ATOMS: Contents

1. Electron Spin (I & II ) (Stern; NP, 1943)
        and the Pauli Exclusion Principle (NP, 1945)
   Electron Configuration and Orbital Diagrams

2. The Aufbau Principle and the Periodic Table
   Orbital Energy Diagram ( Coarse , Fine )
   Filling Order Diagram ( I , II ): mnemonics
   Order of Increasing Orbital Energy (w/ more Symbols )

   Periodic Table

3. Writing Electronic Configuration using Periodic Table(I , II )
   Ga(Z=31) , As(Z=33)
   Electronic Configuration 1 ~ 36

4. Orbital Diagrams of Atoms: Hund's Rule
   Fe (Z=26), H to Na , Ca to Zn

5. Periodic Relationships: Mendeleev's Predictions

6. Some Periodic Properties: Periodic Trend
   Atomic Radius (Table, Figure , Ups Downs ),
   Ionization Energy (Table , Figure ),
   Electron Affinity (Table , Figure)
   ELECTRONEGATIVITY ,
   Summary

7. Brief Descriptions of the Main-Group Elements:
   Overview , Metal vs. Non-Metal,
      Combustion Products from Coal/Woods
   Hydrogen, Group I, Group II, Group III, Group IV,
   Group V, Group VI , Group VII , Group VIII

8. Origin of Elements *: "How Universe and matter were created?"
   A Briefer History of Time: Evolution of Matter
   Formation of Light & Heavy Elements in a Star:
   (Twinkle Twinkle Little Star How I Wonder What You Are ?)
   Evolution of Stars and Explosions of Stars (Supernova)
   Cosmic and Terrestrial Abundance of Elements .
 

Chapter 9 BONDING - GENERAL CONCEPTS

0. Types of Chemical Bonds: Chemical Bonding
   Why Changes (Bonding) Occurs? : Examples
   Formation of a Bond ,

1. Ionic Bonds : Lewis Electron Dot Symbols
   Energy Involved in Ionic Bond ,
   Lattice Energy, Born-Haber Cycle , NaCl Summary

2. Electronic Configuration of Ions :
   Ions of the Main Group Elements
   Transition Metal Ions, Iron Ion

3. Sizes of Ions: LiI , Ionic Radii , Figure

4. Covalent Bond :
   Potential Energy Curve & Orbital Overlap
   Lewis Formulae: Definition , Coordinate Covalent Bond
   Octet Rule, Multiple Bonds

5. Electronegativity ( Table ) and Polar Covalent Bond (Alt.) :
   Bond Polarity and Dipole Moments

6. Lewis Electron-Dot Formula : How to Draw it? (alt. )

7. Delocalized Bonding : Resonance

8. Exceptions to the Octet Rule (alt. ): BF₃

9. Formal Charge and Lewis Formulae: A Question w/ Phosgen
   Formal Charge Method & Bonding Arm Method

10. Properties of Bond : Bond Length and Bond Order

11. Bond Energy (alt. ): Table , Heat and Bond Energy

CH 7 Chemical Bonding: Ionic & Covalent (Whitten et al)

Chapter 10 MOLECULAR GEOMETRY and CHEMICAL BONDING THEORY

1. The Valence-Shell Electron Pair Repulsion (VSEPR) Model
   VSEPR & Molecular Structure,
   VSEPR Theory, Intro : Examples , Analogy , Tetrahedral ,
   How to Predict the Structures (in Steps) , Steric Numbers ,
   How to Predict: Examples (AX ₂ Type) ,
   Geometry (Alternate file ), Summary Table , Structural Examples

2. Dipole Moment (a Vector) and Molecular Geometry
    "How can you tell a molecule is polar or not?"     Symmetry
    H₂O, NH₃, in the Field, Vector Sum,

3. Valence Bond Theory.

4. Description of Multiple Bonding

5. Molecular Orbital Theory

6. Electronic Configuration of Diatomic Molecules

7. Molecular Orbitals and Delocalized Bonding.

Supplemental PowerPoint Lecture Material

             PowerPoint Lecture Slides I (C)

             PowerPoint Lecture Slides II(BLB), 

- FINAL EXAMINATION - COMPREHENSIVE (Chap. 1-10)
A Standard Exam of the American Chemical Society
An Old Sample Final Test (non-Standardized)