JEFFERSON COLLEGE COURSE SYLLABUS CHM112 GENERAL CHEMISTRY II. 5 Credit Hours. Prepared by: Richard A. Pierce. Revised Date: August 2009 by Sean Birke

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1 JEFFERSON COLLEGE COURSE SYLLABUS CHM112 GENERAL CHEMISTRY II 5 Credit Hours Prepared by: Richard A. Pierce Revised Date: August 2009 by Sean Birke Arts & Science Education Dr. Mindy Selsor, Dean

2 CHM112 General Chemistry II I. CATALOGUE DESCRIPTION A. Prerequisite: CHM111 B. 5 Credit Hours C. General Chemistry II applies the principles learned in General Chemistry I to more advanced topics. General Chemistry II covers the stoichiometry of acid-base and oxidation-reduction reactions, chemical thermodynamics, reaction kinetics, chemical equilibrium, and electrochemistry, with special emphasis placed on equilibria in aqueous solutions. Laboratory time is required. (S, Su) II. EXPECTED LEARNING OUTCOMES/ASSESSMENT MEASURES Demonstrate a knowledge of the types of intermolecular forces and their relationship to the properties of matter and to phase changes. Exhibit a detailed knowledge of the structure of the solid state and the bonding in solids. Describe the different types of solutions and the solution process Calculate the concentrations and colligate properties of solutions Understand the three major definitions of acids and bases. Demonstrate a knowledge of colloids and their properties. Relate molecular structure to acid-base strength. Do stoichiometric calculations for acid-base systems using both the mole method and molarity, and the method of equivalent weights and normality Do stoichiometric calculations for redox reactions using both the mole method and molarity, and the method of equivalent weights and normality Understand the laws of thermodynamics and apply them to calculate thermodynamic quantities for chemical reactions. Understand the relationship between thermodynamics and the spontaneity of chemical reactions, and to be able to predict whether a reaction can occur

3 Understand why rates of chemical reactions depend on the conditions under which they occur Explain how to elucidate the mechanism of a chemical reaction. Do kinetic calculations based on the rate law for a chemical reaction Understand the concept of chemical equilibrium and be able to calculate the compositions of equilibrium systems. Understand the relationship between thermodynamics, kinetics, and equilibrium. Understand how to apply equilibrium theory to and do equilibrium calculations for ionic equilibria, i.e., acid-base equilibria, buffer solutions, hydrolysis equilibria, and solubility equilibria. Understand the processes that occur in electrochemical cells. Do electrochemical calculations to predict whether a reaction will occur and to predict the products of the reaction Understand the relationship between equilibrium and the functioning of electrochemical cells III. COURSE OUTLINE WITH UNIT OBJECTIVES A. Liquids, Solids and Phase Changes 1. Polar covalent bonds and dipole moments 2. Intermolecular forces 3. Properties of liquids 4. Phase changes 5. Evaporation, vapor pressure and boiling point 6. Kinds of solids 7. X-ray crystallography 8. Unit cells, the packing of spheres in crystalline solids 9. Structures of some ionic solids 10. Structure of some covalent network solids 11. Phase diagrams B. Solutions and Their Properties 1. Solutions 2. Energy changes and the solution process 3. Units of concentration 4. Some factors affecting solubility

4 5. Physical behavior of solutions: Colligative properties 6. Vapor pressure lowering of solutions: Raoult s Law 7. Boiling point elevation and freezing point depression of solutions 8. Osmosis and osmotic pressure 9. Some uses of colligative properties 10. Fractional distillation of liquid mixtures C. Chemical Kinetics 1. Reaction rates 2. Rate laws and reaction order 3. Experimental determination of a rate law 4. First order reactions a. Integrated rate law b. Half life 5. Second order reactions a. Integrated rate law b. Half life 6. Zeroth order reactions 7. Reaction mechanisms 8. Rate laws for elementary reactions 9. Rate laws for overall reactions 10. Arrhenius Equation a. Temperature and reaction rates b. Applications 11. Catalysis a. Heterogeneous b. Homogeneous D. Chemical Equilibrium 1. The equilibrium state 2. The equilibrium constant K c 3. The equilibrium constant K p 4. Heterogeneous equilibria 5. Using the equilibrium constant 6. Le Chatelier s Principle 7. Factors influencing chemical equilibrium a. Concentration b. Pressure and volume c. Temperature d. Catalyst 8. The connection between chemical equilibrium and kinetics E. Aqueous equilibria: acids and bases 1. Acid Base concepts a. Arrhenius Theory b. Bronsted-Lowry Theory

5 c. Lewis Theory 2. Acid strength and base strength 3. Hydrated protons, hydronium ions 4. Dissociation of water 5. ph scale 6. Measuring ph 7. The ph in solutions of strong acids and bases 8. Equilibria in solutions of week acids 9. Calculating equilibrium concentrations in solutions of weak acids 10. Percent dissociation in solutions of weak acids 11. Polyprotic acids 12. Equilibria in solutions of weak bases 13. Relation between K a and K b 14. Acid base properties of salts 15. Factors that affect acid strength 16. Lewis acids and bases F. Applications of aqueous equilibria 1. Neutralization reactions 2. The common ion effect 3. Buffer solutions 4. The Henderson Hasselbach Equation 5. ph titration curves a. Strong acid strong base titrations b. Weak acid strong base titrations c. Weak base strong acid titrations d. Polyprotic acid strong base titrations 6. Solubility equilibria 7. Measuring K sp and calculating solubility from K sp 8. Factors that affect Solubility 9. Precipitation of ionic compounds 10. Separation of ions by selective precipitation 11. Qualitative analysis G. Thermodynamics: entropy, free energy and equilibrium 1. Spontaneous processes 2. Enthalpy, entropy and spontaneous processes 3. Entropy and probability 4. Entropy and temperature 5. Standard molar entropy and standard entropies or reaction 6. Entropy and the second law of thermodynamics 7. Free energy 8. Standard free energy changes for reactions 9. Standard free energies of formation 10. Free energy changes and reaction mixture composition 11. Free energy and chemical equilibrium

6 H. Electrochemistry 1. Galvanic Cells 2. Shorthand notation for galvanic cells 3. Cell potentials and free energy changes for cell reactions 4. Standard reduction potentials 5. Using standard reduction potentials 6. The Nernst Equation 7. Electrochemical determination of ph 8. Standard cell potentials and equilibrium constants 9. Batteries 10. Fuel cells 11. Corrosion 12. Electrolysis and electrolytic cells 13. Commercial applications of electrolysis 14. Quantitative aspects of electrolysis I. Hydrogen, oxygen and water 1. Hydrogen 2. Isotopes of hydrogen 3. Preparation and uses of hydrogen 4. Reactivity of hydrogen 5. Binary hydrides 6. Oxygen 7. Preparation and uses of oxygen 8. Reactivity of oxygen 9. Oxides 10. Peroxides and superoxides 11. Hydrogen peroxide 12. Ozone 13. Water 14. Reactivity of water 15. Hydrates IV. METHOD OF INSTRUCTION A. Lectures B. Laboratory Work C. Reading Textbook D. Class Discussion

7 V. REQUIRED TEXTBOOK McMurry, John and Fay, Robert. Chemistry, 5 th ed., Prentice Hall, Upper Saddle River, New Jersey, Nelson, John H. and Kemp, Kenneth C. Chemistry, The Central Science Laboratory Experiments, 10 th ed., Prentice Hall, Upper Saddle River, New Jersey, VI. REQUIRED MATERIALS Writing Paper, Pencils, Pens Laboratory Notebook Scientific Calculator VII. SUPPLEMENTAL REFFERENCES None VIII. METHOD OF EVALUATION A. Lecture Examinations B. Laboratory Reports C. Lecture Quizzes D. Grading Scale: A: B: C: D: F: 0-59 IX. ADA STATEMENT Any statement requiring special accommodations should inform the instructor and the Coordinator of Disability Support Services (Library; phone , ext. 169). X. ACADEMIC HONESTY STATEMENT All students are responsible for complying with campus policies as stated in the Student Handbook. See the college website at