The rate at which a chemical reaction occurs can vary depending on a number of factors.  Some reactions occur at a very slow rate, such as the rusting of iron.  Others, like the precipitation of lead (II) iodide, occur very quickly.  Many reaction rates depend on the temperature, the concentrations of the reactants, and the presence of a catalyst.  We are going to focus on the effect of concentration in this lab.

For a general reaction A + B  C, the rate (v) of the reaction is given by the equation

v = k [A]n [B]m

where [A] and [B] are the concentrations of the reactants expressed in Molarity.  The term k is the rate constant for the reaction which must be found experimentally.  The exponents on the concentrations are called the reaction orders of the reactants, and the total of the orders is the reaction order for the overall reaction.  We will be finding the rate law for a reaction between acetone and iodine as follows

C3H6O + I2  C3H5OI + H+ + I-

In this reaction, the rate depends on the concentration of the product H+ as well as the reactants.  Thus, the rate law can be written

v = k [Ac]m [I2]n [H+]p

We wish to find k and the values of the exponents.  In order to experimentally find the reaction rates, we can take advantage of the fact that iodine has a distinct color.  We know that the reaction is done when the iodine is all used up (i.e., when the color is gone).  The experimental rate will simply be

v = D[I2] / t

which is just the change in iodine concentration per unit of time.  We can judge the iodine concentration by watching the color of the sample; when the solution becomes clear we know that the iodine concentration is zero.  Now that we can measure the rate of the reaction directly, we can simply run the reaction several times with different concentrations of reactants, calculate the rates, and compare them to find the reaction orders.

Methodology

1.	Set up a data table as shown below.

2.	Place exactly 10.0 mL of 4.0 M acetone into an Erlenmeyer flask.  Add exactly     10.0 mL of 1.0 M HCl to the flask.  Now add 20.0 mL of distilled water to the flask.

3.	Obtain exactly 10.0 mL of 0.0050 M iodine, but do not add it to the flask yet!  Be careful with iodine; it will stain your hands and clothing.

4.	Be ready to time the reaction.  Now, pour the iodine into the flask and watch the time.  Record the time needed for the solution to become clear.

5.	Repeat the experiment and average the times.

6.	Perform the experiment with 20.0 mL acetone, 10.0 mL HCl, 10.0 mL iodine, and 10.0 mL distilled water.  Repeat and average the results.

7.	Perform the experiment with 10.0 mL acetone, 20.0 mL HCl, 10.0 mL iodine, and 10.0 mL distilled water.  Repeat and average the results.

8.	Perform the experiment with 10.0 mL acetone, 10.0 mL HCl, 20.0 mL iodine, and 10.0 mL distilled water.  Repeat and average the results.

Mixture	Acetone volume (mL)	HCl volume (mL)	Iodine volume (mL)	Water volume (mL)	Time
Trial #1 (s)	Time 2 Trial #2 (s)	Average Time (s)	
A	10.0 	10.0 	10.0 	20.0 				
B	20.0 	10.0 	10.0 	10.0 				
C	10.0 	20.0 	10.0 	10.0 				
D	10.0 	10.0 	20.0 	10.0 				

Analysis

First, we must find the experimental reaction rates for each of the mixtures.  We do this by using the time needed for the iodine to be used up.  Since the final iodine concentration is zero, the equation is

v = [I2] / t

Note that [I2] = 0.0010 M in mixtures A, B, and C, and [I2] = 0.0020 M in mixture D.  Calculate the rate for all four mixtures.  Now that we know the rates, we can plug them into the rate law equations.  Unfortunately we have too many variables, so we solve by elimination.  Consider, for example, dividing the rate equation for mixture A by that for mixture B:

vA       k [Ac]Am [I2]n [H+]p
---  =  -----------------------
vB      k [Ac]Bm [I2]n [H+]p

Notice that the [Ac] values are different (one is 10 mL, the other is 20 mL) but the other concentrations are the same.  So the [I2] term, the [H+] term, and the k all cancel out, leaving us with

				vA / vB  =  [Ac]Am / [Ac]Bm

The easiest way to solve for the exponent m is to express the equation logarithmically.  The equation then becomes

          log (vA / vB )
m  =  ----------------------
          log ([Ac]A / [Ac]B)

Don’t worry if you don’t know what a log is; just use the button on your calculator for now.  Your math teacher will tell you more than you ever wanted to know about logs soon enough!  You can find n by using a ratio built from mixtures A and D, and you can find p with a ratio of mixtures A and C.  Round all exponents to the nearest whole number. 

You can find k by plugging all your rates and exponents into the rate law equation,

v = k [Ac]m [I2]n [H+]p

Do this for all four mixtures, then find the average k.  You now know the rate law for the reaction!

Questions

1.	In step 3, why can't you add the iodine to the flask immediately?

2.	Do your calculated k values agree with each other, or are they scattered all over the place?  What does that tell you about your values for m, n, and p?

3.	What does it mean for a reactant to have a reaction order of zero?

4.	A student accidentally adds too much iodine to reaction B.  How will that affect his results?

5.	The constant k changes with temperature.  As it gets hotter, would you expect k to get larger or smaller?  Explain.