CChieh@UWaterloo.ca

Experiments in Chem 123L

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CHEM 123L - a Laboratory Course

CHEM 123L is a laboratory course associated with the lecture courses CHEM 123 and CHEM 125. However, CHEM 123L is a separate course and students of CHEM 123 from the Mathematics and Engineering faculties need not take it. Close to 800 students take this course in the winter term, and 300 in the spring term every year.

The purposes of these experiments are highlighted here from the instructor's point of view. How the experiments are related to the course material is also shown.

During the Lab period, you are busy handling equipments and chemicals, reading labels, and recording data. You focus on the observation and interpretation of experimental results, and write the lab report. The Lab environment is so complicated that you may be overwhelmed. Thus, you may overlook the things you are supposed to learn in a lab course.

The information provided here can be read before your lab period so that you grasp the major objectives of the experiments. The information is not part of the Laboratory Manual. You can also read it after you have done the experiments to review to see if you have achieved the goals set for the experiments.

In all, nine (9) experiments are Scheduled.

Experiments

  1. Sphere packing
  2. Kinetics - rate law
  3. Kinetics - Activation energy
  4. Standardize solution concentrations
  5. Weak Acids and Bases - titration
  6. Buffer Solutions
  7. Indicators
  8. Qualitative Analysis
  9. Electrochemical Cells

Purposes of Experiments

  1. Sphere packing as models of crystal structures

    Use spheres to build models for the simple cubic (sc), body centred cubic (bcc), face centred cubic (fcc) and hexagonal closed packed (hcp) lattices. These are basic structure types of metals. The spheres are also used to model the structure of salts. Since sodium chloride is one of the common and an important salt; you construct a 3-dimensional structure so that you can describe it well.

    The use of spheres as models for these structures is a brilliant scientific method. Using these spheres you can do the following:

    • Distinguish sc, bcc, fcc, and hcp structure types.
    • Correlate the radii of spheres with edge lengths of the unit cells.
    • Demonstrate the tetrahedral and octahedral sites created by sphere packings.
    • Calculate the radii of spheres that will fit exactly into the octahedral and tetrahedral sites of closest packed spheres.
    • Describe the crystal structure of NaCl
    These are also goals of the lectures on the following topics:
    Intermolecular forces
    Face centered
    Simple cubic
    Tetrahedral sites
    Body centered
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  2. The Effect of Concentration on Reaction Rates
    --- the iodide-persulfate reaction

    In CHEM123/125 lectures on the following topics,

    Chemical kinetics
    Reaction rates and orders
    you have learned how the order of a chemical reaction can be determined, and how the rate law can be derived. In this experiment, you actually determine the order and the rate law of the reaction: (NH4)2S2O8 + 2 KI = I2 + (NH4)2SO4 + K2SO4 or S2O82- + 2 I- = I2 + 2 SO42- You will learn these lab skills: Back to experiment listing.

  3. The effect of temperature and catalyst on the reaction rate
    ---the iodide-persulfate reaction

    Using the same reaction as the previous experiment,

    S2O82- + 2 I- = I2 + 2 SO42- but you allow the reaction to proceed at different temperatures to investigate the temperature effect on the reaction rate. The temperature effect is best described by the
    activation energy, Ea, defined in the Arrhenius equation
    k = A e- Ea / R T Furthermore, you also investigate the effect of a catalyst on the reaction rate. You determine the change in activation energy when the catalyst CuSO4 is added.

    The Arrhenius equation is a model to correlate the rate constant k with temperature T. Many reactions follow this model, and it becomes a theory of chemical kinetics. You can apply the same technique to develop a scientific theories.

    The lab skills you have acquired from Experiment 2 are useful for this experiment. Lots of data are collected, and you are required to carry out extensive calculations from the results. The tables layed out in the Lab manual help you organize the information, and the organization of information is an important skill.

    For the evaluation of Ea, the graphic method is suggested. Again, this is a useful technique, but modern computers may have other techniques to derive the same results. Explore them.

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  4. Standardized Solution Concentrations
    ---acid-base and redox titrations

    Although a standardized solution is provided for your experiment, the ultimate purpose of this experiment is to learn the technique of standardization.

    Precise quantities of pure oxalic acid, H2C2O4, can be weighed and used as the primary standard, unlike NaOH which absorbs moisture from the air. Dissolving a known quantity in a known volume enable you to calculate its concentration, which can be used to standardize concentrations of other bases by the reaction,

    H2C2O4 + 2 OH- = 2 H2O + C2O42-.

    Oxalic acid is also a reducing reagent, and it can be used as a standard reagent to standardize the concentration of oxidizing reagent such as potassium permanganate, especially in an acidic solution. The reaction equation is,

    5 H2C2O4 + 2 MnO4- + 6H+ = 10 CO2 + 2 Mn2+ + 8 H2O. The oxidation and reduction are discussed extensively in the following subjects.
    Oxidation states
    Half reactions
    Balance redox equations

    Standardization is an important skill for any laboratory operation. Standardization of concentrations and quantitative determination often require the titration procedure, which involves the cleaning, rinsing, filling, and reading of burets. Use of indicators and detection of equivalent points are some of the skills you have acquired in this experiment.

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  5. Weak Acids and Bases
    ---weak acid-strong base and strong acid-weak base titrations

    Many compounds can be classified as Acids and Bases Thus, the concept of acid and base is very useful for the study of materials.

    Acids can be divided into Strong Acids & Bases, and Weak acids & bases. Strongacids and bases are completely ionized in their solutions, whereas weak acids and bases are not completely ionized. The equilibrium constants for weak acids and bases are Ka and Kb respectively.

    Water is both a weak acid and a weak base, and its autoionization constant is Kw. Acids and bases are relative terms. For a particular acid or base, it constants have the following relationship:

    Ka Kb = Kw

    The ionizations and equilibria of weak acids and bases make them interesting. Their analysis leads to many types of problems. For example:

    Titration
    Indicators
    Hydrolysis

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  6. Buffer Solutions
    ---pH changes in buffer solutions

    A buffer solution contains a weak acid or base and its salt. The pH of such a solution changes very little when a small amount of acid or base is added to it. The reason is due to the equilibria for the weak acids or bases. For a weak acid, HA, the equilibrium is,

    HA = H+ + A- and the we have
         [H+] [A-]                 [A-]                  [A-]
    Ka = --------;  pKa = pH - log---- or pH = pKa + log-----
           [HA]                   [HA]                  [HA]
    
    Please work out a set of equilibrium equations for a weak base BOH as a practice.

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  7. Indicators
    ---determine Ka values of indicators

    Indicators are weak acids or bases. The principle governing their equilibria is the same as that of weak acids and weak bases. Lets treat an indicator as an acid, HIn. It's equilibrium can be represented by,

    HIn = H+ + In- and then we have
         [H+] [In-]                [In-]                  [In-]
    Ka = ---------; pKa = pH - log----- or pH = pKa + log-----
           [HIn]                  [HIn]                  [HIn]
    
    These are exactly in the same form as those given in Experiment 6 for
    weak acids.

    The module Indicators discusses the subject a step further, and it also has some questions to help you understand the color changes of indicators.

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  8. Qualitative Analysis
    ---identify the presence of metal ions in solutions

    Metal ions in solutions can be identified by their reactions with anions, Cl-, SO42-, CO32-, OH-, S2-, and CrO42- to form insoluble salts.

    Groups of Metal Ions

    The anions mentioned earlier form insoluble salts with various metals. They separate metal ions into these groups:

    Silver, mercury(I), lead form insoluble chlorides.

    Aluminum and chromium(III) form insoluble hydroxides in slightly basic solution.

    Mercury(II), lead, copper, bismuth cadmium, tin, arsenic(III), and arsenic (V) form sulfides insoluble in acid solution. However, iron(II), zinc, cobalt, nickel, and manganese form insoluble sulfides in basic solutions.

    The solubility of metal sulfide S2- is particularly interesting, because the concentration of S2-, [S2-], is controlled by the equilibria

    H2S = H+ + HS-,
         [H+] [HS-]
    K1 = ----------
            [H2S]
    
    and HS- = H+ + S2-,
         [H+] [S2-]
    K2 = ----------
            [HS-]
    

    The equilibrium among H2S, H+, HS-, and S2- are are typical equilibrium of Polyprotic Acids. The pH of a solution determines the concentration of S2-, [S2-]. The solubility of a metal sulfite is governed by Heterogeneous Equilibria. For this type of equilibrium, the equilibrium constant is called the solubility product, Ksp, of Insoluble Salts.

    Barium, strontium, mercury(II), lead, calcium, and silver form sparingly soluble salts with SO42-.

    In this experiment, you also learn to use complexing reagents such as ammonia to dissolve a metal salt. You further explore the temperature effect on solubility. Acids are used to dissolve various solids, and this method is often used to prepare solutions for chemical analysis in industry and in biomedical laboratories.

    When you know the details, the explanation of a simple precipitation can be rather complicated.

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  9. Electrochemical Cells
    ---measure the reduction potentials of electrochemical cells

    In this experiment, you actually construct a Battery of the type:

    Pb | Pb2+ || Cu2+ | Cu Both lead and copper change their Oxidation states in order to provide the energy of a battery. Their tendency to react as Half reactions are related to their chemical potentials. This battery consists of two half cells. The two Half-Cell reactions give rise to a cell potential often referred to as the Cell EMF.

    You apply the Nernst equation to construct a concentration cell.

    Cu | Cu2+ 0.001 M || Cu2+ 0.1 M | Cu The concentration difference can be measured as a electrochemical potential. The changes in [Cu2+] due to the formation of copper complex by adding NH3 cause a change in the electrochemical potential.

    You measure the changes in this experiment, rather exciting eh!

    Since all battery operations involve reduction and oxidation reactions, you might want to review Balance redox equations to refresh your techniques.

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cchieh@uwaterloo.ca