FARMINGDALE STATE COLLEGE DEPARTMENT OF CHEMISTRY

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1 FARMINGDALE STATE COLLEGE DEPARTMENT OF CHEMISTRY COURSE OUTLINE: COURSE TITLE: Prepared by: Dr. Victor Huang September 2016 General Chemistry Principles II COURSE CODE: CHM 153 CREDITS: 4 CONTACT HOURS: Lecture: 3 Laboratory: 3 CATALOG DESCRIPTION: PREREQUISITE: A continuation of General Chemistry Principles I, which includes laboratory. Topics include: solutions and their colligative properties, acids and bases, chemical equilibrium, ionic equilibrium, ph, buffers, titration curves, equilibrium of slightly soluble salts, common ion effect, complex ions and solubility, oxidation and reduction balancing, electrolysis, galvanic/voltaic cells, Nernst equation, Gibbs Free Energy and chemical kinetics. General Chemistry Principles I (CHM-152) or equivalent. IMPORTANT NOTE: REQUIRED FOR: GENERAL EDUCATION: BOTH THEORY AND LABORATORY PARTS OF THIS COURSE MUST BE TAKEN CONCURRENTLY IN ORDER TO RECEIVE CREDIT. Liberal Arts and Sciences, Pre-Health and Life Sciences, Bioscience, Computer Science, Medical Laboratory Technology and Manufacturing Engineering Technology This course satisfies 4 credits of the Natural Sciences competency area of the General Education requirements at Farmingdale State College.

2 2 ELECTIVE FOR: Liberal Arts and Sciences Non-Science majors REQUIRED TEXT: Chemistry Molecular Nature of Matter, 7 th edition, by Jespersen and Hyslop, Wiley Publishing Laboratory Manual for General Chemistry Principles, Part II by Giannotti, and Mark et al. REQUIRED SUPPLIES: Calculator, laboratory coat and safety glasses or goggles. Other items as mandated by instructor.

3 3 FARMINGDALE STATE COLLEGE DEPARTMENT OF CHEMISTRY CHM 153 General Chemistry Principles II Lecture Schedule I. Intermolecular Attractive Forces Types of intermolecular attractive forces, influence of intermolecular attractive forces on physical properties, heating and cooling curves, phase diagrams Section I: At the end of this section, the student should be able to: 1. Describe the four types of intermolecular attractive forces 2. Rank the types of intermolecular attractive forces in order of increasing strength 3. Assign which type of intermolecular force is associated with a compound based on its structure 4. Compare relative magnitude of physical properties based on intermolecular attractive forces inherent in structure 5. Interpret heating and cooling curves, including the thermodynamic parameters associated with each step 6. Interpret phase diagrams based on pressure and temperature II. Properties of Solutions, Mixtures of Substances at the Molecular Level. Heats of solutions, heterogeneous and homogenous mixtures, freezing point depression, boiling point elevation, osmosis, molarity, molality, % solution and mole %. Section II: At the end of this section, the student should be able to: 1. Solve solution concentration problems for the molarity, moles or mass given the appropriate ancillary information. 2. Solve for % solution on a mass as well as volume basis, interconvert between % and Molarity.

4 4 3. Determine the mole fraction of solute and solvent, the extent of vapor pressure lowering on a liquid by the addition of solute. 4. Solve for extent of boiling point elevation as well as freezing point depression, and explain how this occurs. 5. Solve solution stoichiometry problems including limiting reagent type. III: Kinetics The Study of Rates of Reactions. Factors affecting reaction rates, Rate Laws, Integrated Rate Laws, Reaction Rate Theories, Activation Energies, Experimental Rate Laws, Catalysis Reactions. Section III: At the end of this section students should be able to: 1. Know the factors that influence the rate of a reaction occurs. 2. Understand how rates of reactions are expressed and how they might be measured (instantaneous vs average). 3. Understand how the rate of reaction is related to the concentrations of the reactants; understand how to relate changes in concentrations of reactants/products with other reactants/products based on a balanced equation. 4. Write a rate law based on a balanced equation. 5. Use experimental data to determine reactant orders, rate constant, and overall rate law for a reaction. 6. Use integrated rate laws to predict the amount of substrate remaining based on the reaction order and the amount of time elapsed. 7. Understand the concept of half-life and how it is related to the rate constant; predict the amount of substrate left after n halflives have elapsed; predict the number of half-lives required for a substrate to decay to a predetermined amount. 8. Understand the relationship between activation energy and rate constant; compare activation energies to determine the rate-limiting step.

5 5 9. Understand the relationship between rate and temperature; calculate the value of the rate constant based on activation energy and termperature. III. Chemical equilibrium: General Concepts. Dynamic Equilibrium, Kc, Kp, Le Chatelier'sPrinciple, Qc vs Kc. Section III: At the end of this section, the student should be able to: 1. Write the equilibrium expression for heterogeneous equilibrium and solve for K eq. 2. Manipulate equilibrium equations and their corresponding K eq values 3. Apply Le Chatelier s Principle to determine if reactant/product will increase, decrease or remain the same when a stress in the form of volume changes, pressure changes, concentration changes, temperature changes or addition of a catalyst is exerted on the original equilibrium. 4. Calculate the numerical value for K eq starting from equilibrium concentrations 5. Solve equilibrium problems to determine the concentration of all species. 6. Predict the direction of an experimental system towards reactants or products by comparing Qc with Kc 7. Learn when to apply simplifying assumptions 8. When simplifying assumptions fail, to be able to use either a quadratic solution or the method of successive approximations. IV. Acid and Bases-A Second Look. Acid-Base theory (Bronsted-Lowry), Lewis acid theory, ph, poh. Section IV: At the end of this section, the student should be able to: l. Write the definitions for Arrhenius, Bronsted-Lowry and Lewis Acids and Bases. 2. Identify Acid-Conjugate Base and Base-Conjugate Acid Pairs.

6 6 3. Identify the hydronium and hydroxide ions. 4. Understand the inverse relationship between acidity/basicity; hydrogen ion concentration/hydroxide concentration; and ph/poh. 5. Calculate [H + ], [OH - ], ph, and poh of strong acid or strong base solutions V. Equilibria in solutions of Weak Acids and Bases. Ka, Kb, pka, pkb, preparation of buffers, acid-base equilibria. Section V: At the end of this section students should be able to: 1. Understand the inverse relationship between an acid and its conjugate base or a base and its conjugate acid 2. Calculate the value of pka or pkb given Ka or Kb for a weak acid or base. 3. Calculate the value of Kb of a conjugate base from the Ka of an acid; calculate the value of Ka of a conjugate acid from the Kb of a base 4. Use Ka values to compare relative acidity of weak acids (and their conjugate bases); use Kb values to compare relative basicity of weak bases (and their conjugate acids) 5. Predict the ph of weak acid or weak base solutions 6. Predict the ph of solutions of ionic salts that result from acidbase neutralization reactions (hydrolysis of salts). 7. Understand the concept of buffers and buffering capacity. 8. Calculate the ph of a buffered solution after the addition of a strong acid or base. 10. Perform calculations to predict the ph at any of the four general regions of an acid-base titration curve (prior to titration; before neutralization; at equivalence point; beyond equivalence point) for a strong acid-strong base titration and a weak acid-strong base titration.

7 7 VI. Solubility Equilibria. Solubility equilibria of sparingly soluble salts Section VI: At the end of this section students should be able to: VII. Thermodynamics. 1. Write the equilibrium expression for an insoluble salt. 2. Calculate Ksp from solubility data. 3. Calculate the solubility of a salt, given its Ksp value. 4. Understand the common ion effect and how it affects the solubility of some solutes. 5. Predict whether a precipitate will form by comparing Qsp with Ksp Internal Energy, work, heat. Spontaneous and non spontaneous process. Entropy, enthalpy, Second law of thermodynamics, Gibbs free energy, and equilibrium constants. Section VII: At the end of this section students should be able to: 1. Understand what thermodynamics means and to discover the kinds of questions it seeks to answer. 2. Understand how thermodynamics deals with the exchange of energy between a system and its surroundings. 3. Understand what a spontaneous change is and how everything that happens can be traced to some spontaneous change somewhere. 4. Understand what influence energy changes have on the tendency for an event to occur spontaneously. 5. Understand the concept of entropy, S, and to see how and entropy increase favors a spontaneous change. 6. Understand the relationship between the free energy change and the work that is available from the chemical reaction.

8 8 VIII. Electrochemistry 7. Manipulate equations and their corresponding free energy values. Galvanic cells, cell potentials and how it relates to free energies, concentrations in galvanic cells, stoichiometry of electrochemical reactions, practical applications of electrochemistry. Section VIII. At the end of this section students should be able to: 1. Describe galvanic/voltaic cells. 2. Diagram and label a galvanic cell, indicating anode/cathode, anode solution/cathode solution, flow of electrons, flow of ions from salt bridge. 2. Write the half-reactions for the anode and cathode electrodes in a galvanic cell. 3. Write the overall balanced equation corresponding to a galvanic cell. 4. Calculate the cell potential of a galvanic cell. 5. Understand the concept of reduction potential; use reduction potentials to identify oxidizing agents and reducing agents 6. Determine spontaneity of redox process based on cell potential. 7. Relate cell potential to free energy and equilibrium constant.

9 9 FARMINGDALE STATE COLLEGE DEPARTMENT OF CHEMISTRY CHM-153 General Chemistry Principles II LABORATORY EXPERIMENTS Week # Experiment 1 Check In and Safety Lecture 2 Titration of an Unknown Vinegar Solution 3 Titration of an Unknown Chloride Salt 4 Freezing Point Depression 5 Kinetics of Crystal Violet Bleaching 6 Spectrophotometric Determination of an Equilibrium Constant 7 Determination of the Ionization Constant (K a ) Weak Acid 8 Qualitative Analysis of Group I Cations 9 Qualitative Analysis of Group II Cations 10 Qualitative Analysis of Group II Cations (Continued) 11 Qualitative Analysis of Group III Cations. 12 Qualitative Analysis of a General Unknown Sample of Cations. 13 Qualitative Analysis of a General Unknown (Continued) 14 Qualitative Analysis of a General Unknown (Continued) 15 Check Out

10 10 FARMINGDALE STATE COLLEGE DEPARTMENT OF CHEMISTRY CHM-153 General Chemistry Principles II GRADING POLICY Lecture The lecture portion of the course constitutes 75% of the final grade. Four unit examinations are given during the semester. No make-up examination will be given for any missed exam. A comprehensive final is optional for students to take, and can be used to replace a missed or poor examination grade. The final lecture grade will be an average of the four best lecture exam grades (comprehensive final counts the same as a unit exam). Laboratory The laboratory portion of the course constitutes 25 % of the final grade. There are 13 laboratory experiments in total. The lowest laboratory grade will be dropped. The final laboratory grade will be an average of all laboratory report grades. FINAL GRADE= 75% LECTURE + 25% LAB.

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