Reaction Rates - completed labs PDF

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Reaction Rates: Chemical Kinetics of the Iodine Clock Reaction. Introduction Chemicals to your mark, get set, Go! The transformation of a chemical species into a new chemical species can occur in a "blink of an eye" (e.g., the chemical reactions leading to an involuntary action) or proceed "slower than molasses in January" (e.g., weathering of a mountain range). The "kinetics" or rate of a chemical reaction can be influenced by a variety of factors including concentrations of reactants and products, temperature and the presence of catalyst. In this experiment, you will first investigate the kinetics of a classical chemical reaction commonly known as the iodine clock reaction. Specifically, you will: a) investigate the influence of reactant concentration and determine the rate law for the reaction b) investigate the influence of temperature on the reaction rate c) investigate the influence of a catalyst on the rate of a decomposition reaction A. Influence of Concentration on Reaction Rates Reaction rates normally increase with an increase in reactant concentrations. For example, wood (a reactant) burns more rapidly in pure oxygen (a reactant) than in air, which contains only about 20% oxygen. Both the concentration and temperature effects will be studied using the following slow reaction: 3HSO3- + IO3- → 3HSO4- + I -

Eq. 1

In order to study a reaction, it is necessary to follow the rate at which products are formed or reactants are consumed. In this experiment, the rate of consumption of the bisulfite ion, HSO3- will be monitored. This is possible because of the following reactions, which take place rapidly as long as reactants are available: 5 I - + IO3- + 6H+

→ 3I2 + 3H2O

Eq. 2

3I2 + 3H2O + 3HSO3- → 3HSO4- + 6 I - + 6H+

Eq. 3

As soon as HSO3- is depleted, I2 from the reaction in Eq. 2 will not be changed to I- by the reaction in Eq. 3. The I2 will quickly accumulate and form a dark blue complex with starch, which is added as an indicator. Starch +I2 → Blue color

Eq. 4

Thus, the time that elapses between mixing the reactants and the appearance of a blue color in the solution is a measure of the rate at which HSO3- reacts. The shorter the time, the greater the rate. In this part of the experiment, you will determine how the reaction rate varies with the concentration of each reactant, which will enable you to write the Rate Law for the reaction.

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Reaction Rates

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Procedure 1. You should work in pairs for this experiment. Each partner is to record the data, do the calculations, and submit a separate report. 2. Take two clean dry 250 mL beakers to the reagent storage area. Pour about 140 mL of 0.03M potassium iodate (KIO3) solution into one beaker, and about 100 mL of 0.05M sodium bisulfite (NaHSO3) solution into the other beaker. Do not use a graduated cylinder. Use the graduations on your beakers to estimate the volumes. Return the beakers to your work area and label them so you can remember their contents. 3. Refer to Table 1. For each trial, use a 10 mL pipet to accurately measure the indicated volume of 0.03M KIO3 solution from the 250 mL beaker into a clean 125 mL flask. The KIO3 solution is the source of the IO3ion of the reaction in Eq. 1. Use the same pipet to accurately add the indicated volume of distilled water to the same 125 mL flask. 4. Use a second 10 mL pipet to measure the indicated volume of 0.05M NaHSO3 from the 250 mL beaker into a second clean 125 mL flask. The NaHSO3 solution is the source of the HSO3- ion of the reaction in Eq. 1. This solution also contains the starch indicator. NOTE: The contents of the two flasks are to be mixed and a stopwatch or sweep-second-hand watch used to measure the time that elapses between mixing and the appearance of a blue color. 5. Quickly pour the contents of flask # l containing KIO3 and water into flask #2 containing NaHSO3. 6. Begin timing immediately upon mixing. 7. During the first 10-15 seconds, swirl the flask contents to completely mix the reactants. 8. When a blue color first appears, stop the stopwatch or otherwise note the time. 9. Record in Table 1 the Reaction Time, in seconds, for each trial. We will assume that the Reaction Rate (in units of s-1) equals the inverse of the Reaction Time (s). Record the appropriate Reaction Rate in Table 1. _____________________________________________________________________________________ B. Influence of Temperature on Reaction Rates (Continue to work in pairs for this part ) Chemical reactions generally speed up when the temperature is increased. Chemists sometimes use the rule of thumb that a 10°C increase in temperature will double the rate of reaction. The influence of temperature on the rate of the iodine clock reaction will be investigated. Reaction temperatures will be established and maintained by water baths. 1. For each trial listed in Table 2, measure the KIO3 solution and distilled water into a clean flask as you did in Part A. Measure the NaHSO3 solution into a clean 6-inch test tube. Note that the concentrations of the reactants will be the same in each trial, so that any change in the reaction rate is due to a change in temperature.

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Reaction Rates

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2. Before mixing the reactants, bring them close to the same temperature by placing them in a water bath made from your largest beaker and by adjusting to the appropriate temperature. Leave them in the bath for at least 5 minutes before mixing. NOTE: The reactants are to be mixed by pouring the test tube contents into the flask and the reaction timed as in Part A. However, the baths must be set to proper temperatures before the flask and test tube are put in. Read steps 3-5 completely before you proceed. 3. Carry the reactions out at 3 temperatures, room temperature (no bath required), about 10 ° C below room temperature (use cold tap water and ice to adjust the bath temperature), and about 10°C above room temperature (use hot tap water to adjust the bath temperature). 4. In each trial, as soon as the reactants are mixed, place the flask containing the reaction mixture back into the water bath and leave it there until the reaction is completed. 5. After the reaction is completed, insert a thermometer into the reaction mixture. Record the observed temperature in Table 2. 6. Record the Reaction Time, in seconds, and the Reaction Rate (s-1) for each reaction in Table 2. _____________________________________________________________________________________ C. Influence of Catalysts on Reaction Rate Some substances, known as catalysts, have the ability to speed up chemical reactions without being consumed in the process. Catalysts operate in a variety of different ways. Some provide an alternate, faster pathway for a reaction by forming intermediate compounds with one or more of the reactants (homogeneous catalysis). Others provide surfaces on which the reactants adsorb and react more easily (heterogeneous catalysis). The catalytic exhaust converter of post-1975 automobiles is an example of a heterogeneous surface catalyst. Whatever mechanism is involved, catalysts are very useful and important because of their rate-enhancing abilities. The effort of catalysts will be studied as a part of this experiment by observing the rate of decomposition of a solution of hydrogen peroxide: 2H2O2 → 2H2O + O2

Eq. 5

This reaction can be followed easily by noting the rate of evolution of gaseous oxygen bubbles from the solution. Procedure 1. The remainder of the experiment should be performed on an individual basis. Be sure that all glassware is exceptionally clean. Otherwise, catalysis by impurities may give false results. 2. Use a clean graduated cylinder and obtain about 15 mL of 3 % hydrogen peroxide. 3. Divide the hydrogen peroxide by pouring approximately equal amounts into 3 small clean test tubes.

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Reaction Rates

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4. Test the catalytic activity of the substances listed in Table 3 by adding 5-6 drops of solutions or about 0.l g of solids to separate test tubes containing peroxide. 5. Mix well, then observe each mixture for 3-4 minutes. 6. Record your observations in Table 3 in terms of rate of evolution of oxygen gas (rapid; moderate; slow; none). CALCULATIONS A. Influence of Concentration on Reaction Rates: Determination of the Rate Law A Rate Law is an equation that gives the rate of a reaction as a function of the concentrations of reactants. The general form of the Rate Law for this reaction is given by

where the exponents x and y represent the order of the reaction in HSO3- and IO3-, respectively. k is called a rate constant, and is constant for a given reaction at a given temperature. The order of a reactant can be determined by isolating that reactant’s effect on the reaction rate. This is done by measuring the reaction rate for two different experiments in which only the concentration of the reactant of interest is changed. Referring to Table 1, for example, the difference in reaction rates in the first two trials must be due to the difference in HSO3- concentration, since the concentration of IO3- is the same in both trials. The ratio of the rates for the two experiments can be expressed in terms of the rate law as

But, since k and [IO3-] are the same for trials 1 and 2, they cancel out, such that the ratio is now given by

We can now solve for the order in HSO3- (which is x), since all other values in the above equation are known. The order in IO3- can be determined similarly by comparing the appropriate two trials.

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Reaction Rates

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Data and Report A. Influence of Concentration on Reaction Rates Table 1: Trial Volume KIO3 (Flask #1)

( Note: Total volume is always 60.0 mL) Volume Volume Initial Concentration water NaHSO3 (right after mixing)/M (Flask #1) [KIO3] [NaHSO3] (Flask #2)

1

20.0 mL

30.0 mL

10.0 mL

2

20.0 mL

20.0 mL

20.0 mL

3

30.0 mL

10.0 mL

20.0 mL

1.

What is the order of this reaction in IO3- ? Show your calculations.

2.

What is the order of this reaction in HSO3- ? Show your calculations.

Reaction Time/s

Reaction Rate/s-1

3. Calculate the value of the rate constant, k, for each of the three trials. Report the mean value of k and its standard deviation. Be sure to report the mean value of k using the correct number of significant figures.

4.

Write the actual rate law for this reaction.

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Reaction Rates

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B. Influence of Temperature on Reaction Rates Table 2: Volume HSO3- Temperature Trial Volume Volume KIO3 (flask) water (flask) (test tube) 1 20.00 mL 30.00 mL 10.00 mL 2 20.00 mL 30.00 mL 10.00 mL 3 20.00 mL 30.00 mL 10.00 mL

Reaction Time

Reaction Rate

1. By what factor do you observe the reaction rate to increase for each 10oC temperature increase? Is this consistent with the ‘Rule of Thumb’ given earlier in the write-up?

C. Influence of Catalysts on Reaction Rates Table 3: Substance Tested Powdered MnO2(s) 0.1 M FeCl3 (aq) 0.1 M NaCl (aq)

Rate of oxygen evolution

1. The two solutions tested are solutions of ionic compounds. List all ions present in the two solutions, and identify each ion as being catalytic or non-catalytic.

2. As we have seen, individual ions in solution can have catalytic power. For many solid catalysts, however, catalytic power is not due to the individual ions or molecules that make-up the solid, but rather to the unusual chemical properties of surface of the solid. This is true of the MnO2 solid that you tested today. How then, would you expect the catalytic activity of 0.1 g of the powder form you used today to compare to that of a single 0.1g sphere of the same material? Explain your answer to receive credit.

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Reaction Rates

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