1. To determine the rate law for a chemical reaction.
2. To use graphical techniques in the analysis of experimental data.
[This laboratory exercise was adapted from one in Beran and Brady' s Laboratory Manual for General Chemistry, John Wiley & Sons, 1982]
The rate of a chemical reaction can be measured in terms of the rate of disappearance of one of the reactants or in terms of the rate of the appearance of one of the products. For example in the hypothetical reaction, A + 2B ----> AB_{2} the rate could be measured in terms of the disappearance of either reactant A or B, or in terms of the appearance of the product AB_{2}. The rate is nearly always proportional to each reactant' s concentration raised to some power. For the example above this means we would write the following relationship: rate=k*[A]^{p}*[B]^{q}. This expression is called the rate equation for the reaction. The value of k, called the rate constant, is unique for a specific reaction and is dependent only on the temperature. A study of the kinetics of any reaction involves determining the values of k, p, and q.
In this experiment you will determine the rate equation for the reaction of potassium persulfate, K_{2}S_{2}O_{8}, with potassium iodide. The reaction is: S_{2}O_{8}^{2-} + 2I^{1-} ----> 2SO_{4}^{2-} + I_{2} .
The rate equation for this reaction will be: rate=k*[S_{2}O_{8}^{2-}]^{p} *[I^{1-}]^{q}.
In logarithmic form this equation becomes: log(rate)=log(k) + p*log[S_{2}O_{8}^{2-}] + q*log[I^{1-}].
In order to determine the value of p you will measure the rate of a series of reactions for a constant iodide ion concentration as the persulfate ion concentration is changed and plot log(rate) vs. log[S_{2}O_{8}^{2-}]. The value of p is the slope of this plot because at a constant iodide ion concentration the logarithmic form of the rate equation reduces to:
log(rate)= p*log[S_{2}O_{8}^{2-}] + C ; where C=k+q*log[I^{1-}].
As can be seen this equation now takes the form of the algebraic " slope intercept equation" (Y=mx + b) where p and m are equivalent. The value of q can be determined in a similar manner.
The reaction rate constant (k) will be determined by substituting the values of p, and q back into the rate equation, calculating k for each trial and averaging the results.
The rate of the reaction will be the time it takes to produce a given quantity, 0.1 mmols, of iodine. To aid in determining when 0.1 mmols of iodine has been produced we will react the iodine produce with sodium thiosulfate, Na_{2}S_{2}O_{3}. This reaction is:
2S_{2}O_{3}^{2-} + I_{2} ----> 2I^{1-} + S_{4}O_{6}^{2-}.
The reaction of Na_{2}S_{2}O_{3} with I_{2} is many orders of magnitude faster than the reaction we are studying so its presence will have no effect on our investigation. Exactly 0.2 mmols of Na_{2}S_{2}O_{3} will be added to the reaction mixture for each trial. Therefore, excess iodine will begin to accumulate, and react with a starch indicator, only after all of the sodium thiosulfate has reacted. Since 0.2 mmols Na_{2}S_{2}O_{3} will consume 0.1 mmols of I_{2}, the starch indicator will turn blue when 0.1 mmols of iodine have been produced by the reaction we are studying. The rate of the reaction can be found by dividing 0.1 mmols iodine by the time in seconds it takes for the blue color to appear.
The table below summarizes the preparation of the test solutions. Measure the volumes of to ±0.05 mL using either a pipet or a buret. All volumes are in mL.
Test Soln |
Solution A | Solution B | |||
---|---|---|---|---|---|
Water | 0.3M KI | starch | 0.02M Na_{2}S_{2}O_{3} | 0.1M K_{2}S_{2}O_{8} | |
1 | 148 | 10 | 2 | 10 | 30 |
2 | 138 | 20 | 2 | 10 | 30 |
3 | 128 | 30 | 2 | 10 | 30 |
4 | 108 | 30 | 2 | 10 | 50 |
5 | 88 | 30 | 2 | 10 | 70 |
1. Prepare solution A in a 250 mL beaker or 250 mL Erlenmeyer flask. Stir the solution thoroughly and record its temperature.
2. The reaction begins when solution B is poured into solution A, therefore, be prepared to start timing the reaction in seconds with a watch or clock. Place the reaction vessel on a white sheet of paper so that the color change is more easily detected.
1. Determine the initial concentration, after mixing, of both persulfate and iodide ions for all trials.
2. Determine the rate for each trial in mmols iodine/sec.
3. Using trials 1,2, and 3 plot log rate vs. log([I^{1-}]). Determine q from this plot.
4. Using trials 3, 4, and 5 plot log rate vs. log([S_{2}O_{8}^{2-}]). Determine p from this plot.
5. Using the p and q determined in steps 3 and 4 and the measured rate, substitute into the rate equation for each trial and determine k. Calculate the average k.
1. State the effect of each of the following changes on the reaction rate in this experiment. Explain your answers.
a) substituting 0.001 M acetic acid for the distilled water in solution A.
b) increasing the concentration of persulfate.
c) increasing the concentration of KI.
d) an increase in the Na_{2}S_{2}O_{3} concentration.
e) increasing the concentration of the starch solution.
f) increasing the temperature of the reactants.
2. Two test reactions are the minimum required to obtain values of p and q in this experiment. Explain the advantage that additional test reactions have on determining p and q as you did in this experiment.
3. What would appear in solution if the Na_{2}S_{2}O_{3} solution were omitted?
4. Write the rate law for the reaction investigated in this experiment with the determined values of p, q, and k.
5. One of the errors in this experiment is in the way we calculated the rate of each trial (dividing 0.1 mmols of iodine by the time in seconds). This is known as the average rate. What kind of rate would have to be calculated for each trial to improve the procedure? Explain.
6. Summarize all calculations in a nice neat table.
Results from the Chemistry II class of Lapeer East High School.
Questions? Comments?? |