Chemical Energetics and Dynamics

Curriculum guideline

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Course
CHEM 1210
Discontinued
No
Course code
CHEM 1210
Descriptive
Chemical Energetics and Dynamics
Department
Chemistry
Faculty
Science & Technology
Credits
5.00
Start date
End term
201930
PLAR
No
Semester length
15 weeks
Max class size
36
Contact hours
Lecture: 4 hours per week Laboratory: 3 hours per week
Method(s) of instruction
Lecture
Lab
Learning activities

The course will be presented using lectures, problem sessions and class discussions.  Films and other audio-visual aids as well as programmed material will be used where appropriate.  Problems will be assigned on a regular basis.  The laboratory course will be used to illustrate the practical aspects of the course material.  Close coordination will be maintained between laboratory and classroom work whenever possible.  This will be accomplished by discussing laboratory experiments in class and when necessary, by using the lab period for problem solving.

Course description
Topics studied will include liquids, solids, solutions, electrochemistry, the laws of thermodynamics, equilibrium, acids and bases and chemical kinetics. A practical laboratory component (requiring good manual dexterity) is an integral part of the course.
Course content

Liquids and Solids

Phases:  vaporization; boiling, vapour pressure, Clausius-Clapeyron equation, other phase changes, phase diagrams (one component) and associated concepts; crystalline solid types, X-ray diffraction, close packing of spheres model, hexagonal and cubic lattices, associated calculations; lattice energies (enthalpies) and Born-Haber cycles.

Solutions

(Review: types of solutions, solution concentrations) The solution process and associated energetics, Henry’s Law, Raoult’s Law (one and two volatile components), fractional distillation.

Electrochemistry

Concentration effects and Nernst equation, relationship between Ecell and ?G/K, electrolysis, quantitative electrolysis. (Review: sufficient review of electrochemical cells and standard reduction potentials to effectively cover the section 3 topics above)

Chemical Kinetics

(Review: basic factors affecting reaction rates) concept and definitions of chemical reaction rates, differential rate law and rate constant, integrated rate laws for zero, first and simple second order reactions, half-life; collision theory and activation energy, reaction profile diagrams, mechanism and rate equations, steady-state approximation, SN1 and SN2 reaction mechanisms, homo- and heterogeneous catalysis.

Equilibrium

(Review: basic principles of chemical equilibrium, equilibrium constant (K) and expressions, magnitude of K, basic Le Chatelier’s principle) Kc versus Kp, reaction quotient, homogeneous (Kp focus) and heterogenous equilibrium calculations (including approximations), detailed examples of Le Chatelier’s principle.

Thermodynamics

(Review: basic concepts of thermochemistry, simple heat capacity problems) First Law of Thermodynamics, calorimetry (constant pressure and volume), enthalpy, Hess’s Law, standard enthalpies of formation, entropy, standard molar entropies, Third Law, Second Law and derivation of Gibbs free energy, standard free energies of formation, free energy and spontaneity, relationship between free energy and equilibrium, thermodynamic equilibrium constants, temperature dependence of equilibrium constants.

Acids and Bases

(Review: Arrhenius and Bronsted-Lowry theory, auto-ionization of water and and Kw, pH, strong/weak acids and bases, Ka, Kb, qualitative hydrolysis of salts) quantitative hydrolysis of salts, polyprotic acids, common ion effect, buffer solutions, titration curves (strong and weak acids/bases), indicators, solubility product (Ksp).

Laboratory Course Content

  1. Synthesis of a Coordination Compound
  2. Oxidation/Reduction Analysis
  3. Spectrophotometric determinations
  4. Kinetics
  5. Thermochemistry
  6. Equilibria
  7. pH and Indicators
  8. Electrochemistry
  9. Acids and bases

NOTE: The student must be able to physically accomplish the various tasks involved in the experiments above.  This includes, for example, sufficient manual dexterity for accurate manipulation and use of volumetric glassware.

Learning outcomes
  1. Define or explain any of the chemical terms used in the course (i.e. anode, state function).
  2. Draw the unit cells for the three cubic lattices.
  3. Given the unit cell of an ionic compound, predict the simplest formula.
  4. Describe the experimental method for obtaining the dimensions of the unit cell.
  5. Explain the differences between cubic and hexagonal closest packing of spheres.
  6. Describe (or draw) the cubic unit cells and ionic crystal structures.
  7. Describe the method of calculating lattice energies using the Born-Haber cycle.
  8. Solve problems of the following types, given a list of selected equations:
    • Determination of the amount of material produced in an electrolytic cell
    • Calculation of the e.m.f of a galvanic cell
    • Calculation of ?G from electrochemical data
    • Calculations involving use of the First Law of Thermodynamics
    • Enthalpy changes in a chemical or physical process
    • Hess’s Law
    • Calculation of ?S from absolute entropies
    • Calculation of ?G for a chemical reaction
    • Calculation of K from ?G°
    • Equilibria in gaseous systems
    • Equilibria in aqueous acid-base systems (pH, weak acids, hydrolysis, buffers)
    • Order, rate constant and activation energy of a chemical reaction
    • Amounts of material involved in redox reactions
  9. Determine the mass of a substance involved in a redox reaction.
  10. State Faraday’s Law of Electrolysis.
  11. Determine whether chemical reactions will occur spontaneously under standard conditions, given a table of standard electrode potentials.
  12. Using a table of standard electrode potentials, compare the relative strengths of oxidizing agents or reducing agents.
  13. Distinguish between various types of heats of reaction and write the corresponding chemical equation.
  14. Interpret the signs of enthalpy changes.
  15. Describe both qualitatively and quantitatively the contributions of ?H and ?S to reaction spontaneity.
  16. Predict the sign of ?S for various chemical and physical processes.
  17. Interpret equilibrium in terms of the thermodynamic driving forces.
  18. Write the chemical equation for the equilibrium involving weak acids and bases in aqueous solution.
  19. Classify various aqueous salt solutions as acidic, basic or neutral and write the corresponding equation.
  20. Explain how an acid-base indicator works and choose suitable indicators for various acid-base reactions.
  21. Verify that the proposed mechanism of a chemical reaction is consistent with the experimentally determined rate law.

Laboratory Objectives

The student will be able to:

  1. Give the name and describe the use of some of the more common laboratory equipment.
  2. Accurately perform standard laboratory techniques using the accepted methods, such as titration, weighing, pipetting.
  3. Give the random and systematic errors inherent in each of the common quantitative techniques which are used in the laboratory.
  4. Given an experimental problem, state the series of steps and the accepted techniques required to solve that problem in the laboratory.
  5. Write a report based on observations and data obtained in the laboratory using a standard report format.
  6. Given a set of experimental data or using data obtained in the laboratory, apply the appropriate mathematical techniques (e.g. graphical analysis, solution of equations, etc.) necessary to obtain a numerical result.
  7. Using the data, observations or results of an experiment, determine the relationship between experimental variables.
  8. Analyze the overall laboratory experiment with respect to errors inherent in the method or techniques.
  9. Give the theory upon which the experiment is based.
Means of assessment
  1. Laboratory work (30%)
    • Laboratory reports:  14%
    • Laboratory practical exam: 8%
    • Qualitative results of experiments performed on unknown samples will be graded: 8%
  2. Examinations (70%)
    • A final comprehensive examination during the exam period:  30%
    • A minimum of two in class tests will be given throughout the semester:  30%
    • Any or all of the following evaluations, at the discretion of the instructor: problem assignments, quizzes, class participation [5% maximum] (10% in total)

Note:

A student who misses three or more laboratory experiments will earn a maximum P grade.

A student who achieves less than 50% in either the lecture or laboratory portion of the course will earn a maximum P grade.

Textbook materials

Textbooks and Materials to be Purchased by Students

R.H. Petrucci & W.S. Harwood and F.G. Herring:  General Chemistry, Douglas College Chemistry 1110/1210, Pearson Custom Publishing, 3rd Ed., Boston 2007.

Douglas College Laboratory Manual Chemistry 1210

Prerequisites

CHEM 1110, C or better

Corequisites

Courses listed here must be completed either prior to or simultaneously with this course:

  • No corequisite courses
Equivalencies

Courses listed here are equivalent to this course and cannot be taken for further credit:

  • No equivalency courses
Which prerequisite

CHEM 2310, CHEM 2315, CHEM 2410