Course Outline: General Chemistry Principles II (CHM153)

Credits: 4
Contact Hours: Lecture: 3
Lab: 3

NOTE on Laboratory: Both Lecture and Laboratory must be taken simultaneously; separate grades will not be given for either. Students must pass the laboratory section to receive a passing grade in the entire course.

Semesters Offered: Fall, Spring, & Summer

Prerequisites: General Chemistry Principles I (CHM 152) or equivalent.

Catalog Course Description:

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.

Required Course for:

Bioscience; Medical Laboratory Technology; Certificate for Health Professions

Elective Course for: 

General Education:

Course Texts: 

 This course satisfies 4 credits of the Natural Sciences competency area of the General Education requirements at Farmingdale State College.

Chemistry: The Molecular Nature of Matter (Jespersen and Hyslop, 7th Edition, Wiley)

CHM 152 Laboratory Manual for General Chemistry Principles (Giannotti, Mark, et al., FSC Chemistry Dept.)

Other Required Course Materials

Calculator, laboratory coat and safety goggles.

Course Learning Objectives:

I. Intermolecular Attractive Forces

Types of intermolecular attractive forces, influence of intermolecular attractive forces on physical properties, heating and cooling curves, phase diagrams

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 %.

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.

        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

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 half-lives 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.

        9. Understand the relationship between rate and temperature; calculate the value of the rate constant based on activation energy and termperature.

IV. Chemical equilibrium: General Concepts

Dynamic Equilibrium, Kc, Kp, Le Chatelier'sPrinciple, Qc vs Kc.

At the end of this section, the student should be able to:

  1. Write the equilibrium expression for heterogeneous equilibrium and solve for Keq.
  2. Manipulate equilibrium equations and their corresponding Keq 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 Keq 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.

V. Acid and Bases-A Second Look

Acid-Base theory (Bronsted-Lowry), Lewis acid theory, pH, pOH. 

At the end of this section, the student should be able to:

  1. Write the definitions for Arrhenius, Bronsted-Lowry and Lewis Acids and Bases.
  2. Identify Acid-Conjugate Base and Base-Conjugate Acid Pairs.
  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

VI. Equilibria in solutions of Weak Acids and Bases

Ka, Kb, pKa, pKb, preparation of buffers, acid-base equilibria.

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 acid-base 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.
  9. 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.

VII. Solubility Equilibria

Solubility equilibria of sparingly soluble salts

At the end of this section students should be able to:

  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

VIII. Thermodynamics

Internal energy, work, heat; Spontaneous and nonspontaneous processes;  Entropy, enthalpy, Second Law of Thermodynamics, Gibbs free energy, and equilibrium constants.

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.
  7. Manipulate equations and their corresponding free energy values.

IX. Electrochemistry

Galvanic cells, cell potentials and how it relates to free energies, concentrations in galvanic cells, stoichiometry of electrochemical reactions, practical applications of electrochemistry.

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.
  3. Write the half-reactions for the anode and cathode electrodes in a galvanic cell.
  4. Write the overall balanced equation corresponding to a galvanic cell.
  5. Calculate the cell potential of a galvanic cell. 
  6. Understand the concept of reduction potential; use reduction potentials to identify oxidizing agents and reducing agents
  7. Determine spontaneity of redox process based on cell potential.
  8. Relate cell potential to free energy and equilibrium constant.

Laboratory Schedule

Experiment                                     Title                                                                                   Pages

1                                                     Check-In & Safety Lecture

2                                                     Titration of Unknown Vinegar Solution                                  1-7

3                                                     Titration of Unknown Chloride Solution                                 8-15

4                                                     Freezing Point Depression                                                   16-26

5                                                     Kinetics of Crystal Violet Bleaching                                      27-45

6                                                     Determination of Equilibrium Constant of                             46-57

                                                                [FeSCN]2+ by Spectrophotometry

7                                                     Determination of Ionization Constant (Ka) of                        58-73

                                                                a Weak Acid

8                                                     Qualitative Analysis 1 (Group I Cations)                               74-81

9                                                     Qualitative Analysis 2 (Group II Cations)                              82-94

10                                                   Qualitative Analysis 3 (Group III Cations)                             95-105

11                                                   General Unknown                                                                  106-116

12                                                   General Unknown, cont.

13                                                   General Unknown, cont.

14                                                   Checkout