CHM 2046 - General Chemistry II

College of Natural Sciences

Credit(s): 3
Contact Hours: 47
Effective Term Fall 2018 (550)

Requisites

Prerequisite CHM 2045 with a minimum grade of C and
Prerequisite CHM 2045L with a minimum grade of C and
Pre- or Co-requisite CHM 2046L with a minimum grade of C

Course Description

This course is a continuation of General Chemistry I and extends the study of chemical principles in solutions, kinetics, gaseous and solution equilibria, acid-base reactions, thermodynamics, oxidation-reduction, electrochemistry, and nuclear chemistry.

Learning Outcomes and Objectives

  1. The student will define and apply concentration terms used in solution chemistry and the colligative properties of solutions by:
    1. using Henry's law, given data of one set of conditions, determine the concentration of a gas in a solution.
    2. using Rault's law, calculate the vapor pressure of a solution
    3. performing calculations related to molarity, molality, mass percent, and mole fraction
    4. performing calculations and applying concepts related to colligative properties: freezing point depression and boiling point elevation
  2. The student will apply the principles of chemical kinetics by:
    1. determining the order of a reaction, given the rate as a function of concentration of reactants.
    2. determining the rate of a reaction, given the concentration of reactant as a function of time.
    3. writing a rate expression for the reaction, and calculating the rate constant given the rate at a known concentration when given the order of a reaction.
    4. using rate equations to determine original concentrations and the rate constants.
    5. using the rate equations to determine the time required for the concentration of reactant to drop to a particular value, given the rate constant and the original concentration. Also determining the initial concentration given the concentration at some particular time.
    6. calculating the other quantity when given either the half-life or rate constant for a first order reaction.
    7. describing and assessing energy diagrams showing energy of activation and enthalpy change and describing the effect of catalysis.
    8. listing and describing the three factors which effect rates of reaction according to collision theory.
    9. using the Arrhenius equation to obtain the rate constant at T2 given its value and T1 and the energy of activation.
    10. using the Arrhenius equation to obtain the activation energy given rate constants at two different temperatures.
  3. The student will describe the nature of gas phase equilibrium systems by:
    1. writing the corresponding expressions for Kp and KC when given a balanced equation for a reaction involving gases.
    2. interpreting the magnitude of KC in relation to the extent of forward and reverse reactions.
    3. using a given equation, calculate the numerical value of KC, and know the equilibrium concentrations of all species.
    4. calculating the numerical value of KC, given the original concentrations of all species and the equilibrium concentration of one species for a given chemical equation.
    5. predicting the direction in which a chemical system will move to reach equilibrium when given the value of KC.
    6. predicting the equilibrium concentration of one species, knowing the concentrations of all other species at equilibrium when given the value of KC.
    7. predicting the equilibrium concentrations of all species, given their initial concentrations when given the value of KC.
    8. predicting the effect of a change in number of moles, volume, or temperature upon the position of an equilibrium by using Le Chatelier's Principle.
  4. The student will explain the nature of aqueous solution systems by:
    1. predicting whether it will be an electrolyte or a nonelectrolyte in aqueous solution when given the formula for a substance.
    2. predicting the relative solubilities of different solutes in water.
    3. predicting the effect on solubility of a change in temperature or pressure.
    4. writing balanced net ionic equations for the formation of a solution, and for the formation of precipitates.
    5. using an equation for a precipitation reaction, relating the amounts of reactants and products.
  5. The student will apply acid-base chemical principles by:
    1. giving the following: [H+], [OH-], pH, or pOH calculating any of the others.
    2. writing equations for the dissociation of strong acids or strong bases.
    3. writing equations for the dissociation equilibria of weak acids and weak bases in aqueous solutions.
    4. predicting whether a given substance will give an acidic, basic, or neutral aqueous solution, and writing an equation for the solution system.
    5. writing equations for the reactions of acids and bases, and describing the solutions that result as acidic, basic, or neutral.
    6. using titration data for an acid-base reaction to determine: the concentration of an acid or a base in aqueous solutions, and molecular mass of an acid or a base.
    7. selecting an acid-base indicator appropriate for a given acid-base titration.
    8. classifying any given species in a reaction as an acid or base, according to the Arrhenius, Bronsted and Lowry, or Lewis models and indicating the conjugate acid-base pairs.
  6. The student will explain the nature of aqueous ionic equilibrium systems by:
    1. writing an equilibrium constant expression for Ksp, for the solution of as lightly soluble ionic substance; Ka for the dissociation of a weak acid in water solution; Kb, for the reaction of a weak base with water; and Kf, for the formation of a complex ion.
    2. using equilibrium constant expressions to determine equilibrium concentrations and equilibrium constants associated with ionic equilibria.
    3. predicting the formation of precipitates or solubility using solubility product, KSP data.
    4. using KA, KB, and KW as well as appropriate equilibrium constant expressions to predict equilibrium concentrations in acid-base systems.
    5. using acid and base dissociation constants to determine pH and concentrations in buffer solutions.
    6. determining the charge on the central metal atom when given the formula of a complex ion or coordination compound.
    7. calculating the equilibrium constant for a reaction, given the equilibrium constant for the reverse reaction, and/or equilibrium constants for two or more other, related reactions. Stating the law of multiple equilibria.
  7. The student will apply the principles of chemical thermodynamics by:
    1. calculating the standard entropy change for a reaction when given the standard molar entropies of reactants and products.
    2. giving the enthalpy change and the standard entropy change for a reaction, calculate the standard free energy change at 298K and at any other temperature.
    3. calculating the temperature at which equilibrium will exist at 1 atmosphere when given the enthalpy change and the standard entropy change for a reaction.
    4. quantitatively relating the standard free energy change and the E0 for a given reaction at 298K.
    5. quantitatively relating the standard free energy change and the equilibrium constant K for a reaction in an aqueous system.
    6. using the laws of thermochemistry to calculations involving standard entropy change, standard free energy change, and enthalpy change.
  8. The student will apply the principles of redox reactions and electrochemistry by:
    1. determining the oxidation number of each atom in a molecule or an ion when given the molecular or ionic formula.
    2. balancing redox reactions using the half-equation method.
    3. labeling the oxidizing and the reducing agents and the species being oxidized and reduced in a balanced oxidation-reduction reaction.
    4. using standard voltages to: decide whether or not a given redox reaction will occur at standard concentration and pressure at 298K.
    5. using a given redox reaction, write the expression for the Nernst equation and use the equation to calculate: the voltage E of a cell, given E0, and the concentrations of all other species.
    6. using a reaction, write the expression for the Nernst equation and then use it to calculate the cell voltage, given the standard voltage and concentrations of all species.
    7. calculating the concentration of one of the reactant species when given the balanced equation for a redox reaction and titration data for the reaction.
  9. The student will explain principles of radioactive decay by:
    1. identifying and describing the types of nuclear changes in an atom.
    2. quantitatively relating nuclear changes with atomic mass changes.
    3. predicting nuclear stability from atomic number and atomic mass.
    4. calculating the amount of a decaying substance remaining after a specified length of time.
    5. calculating the mass and energy changes for nuclear reactions.
    6. describing nuclear fusion and nuclear fission.

Criteria Performance Standard

Upon successful completion of the course the student will, with a minimum of 70% accuracy, demonstrate mastery of each of the above stated objectives through classroom measures developed by individual course instructors.

History of Changes

Revised 7/25/83 Revised 8/84 DBT 2/86 Effective Session 19861 DBT 7/16/87 Effective Session 19871 DBT 5/15/90 Effective Session 19901 3 YR C&I Review 9394 3 YR C&I Review 9899 C&I 11/9/99, DBT 12/14/99 Effective 19992. 3 Year Review 2003. 3 Year Review 2007: C&I 11/13/07, BOT 12/17/07, Effective 20072(0390). C&I Approval: 10/12/2010, BOT Approval: 11/16/2010, Effective Term: Spring 2011 (435).
C&I Approval: 02/09/2018, BOT Approval: 04/17/2018, Effective Term: Fall 2018 (550)

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