(Adapted from Williamson, Little, Microscale Experiments for General Chemistry, Houghton-Miflin, 1997. )
1 lab period; work in groups. Complete the Preparation page before laboratory.
Goals
Background
Chemical kinetics is the study of the the rates of chemical reactions, reflected in the time-dependence of the concentrations of reactants and products. The rate of a chemical reaction is generally expressed as the change in the concentration of a reactant or product per unit time. For reaction (1), then, showing nitrogen and hydrogen reacting to form ammonia, rate may be expressed in three ways, as shown.
Because by convention chemists treat reaction rate as a positive quantity, negative signs are placed in front of the concentration changes for the reactants, because reactants are used up as reaction proceeds and therefore experience negative concentration changes. Ammonia is formed as reaction proceeds, so its concentration change is positive. Other than this matter of sign, however, there is another minor problem that must be addressed. As a result of the reaction stoichiometry, the changes in concentration of the three substances in a given time period are not the same, and unless we correct for this somehow, the rate will be different depending on which of the substances we choose to express it. This problem is readily corrected by dividing the concentration change for a species by the stoichiometric coefficient for the species. Thus,
For essentially all chemical reactions, the rate depends upon the same factors. These are
Usually the rate depends on the concentrations of reactants (and products) according to a simple mathematical form. This is shown for generic reaction 2) in 3):
Here m and n are generally NOT the same as the stoichiometric coefficients, though they are usually integers, often 1 or 2. The number m is called the order of the reaction for reactant A, and n is the order for reactant B. The total reaction order is m+n.
The rate constant, k, implicitly contains within its magnitude the dependence of rate on the chemical nature of the participants and the temperature dependence of the rate. The latter can be independently measured by running the reaction at a series of temperatures. The minimum goal of a kinetics study is to determine the rate law (that is, on which concentrations does rate depend, and what are the orders in these concentrations) and the value of the rate constant, k, at a particular known temperature. Frequently the study is extended to other temperatures to demonstrate the manner in which k varies as T increases. Almost invariably, rate constants (hence rates) increase as T goes up.
There are a number of experimental approaches to determining the rate law. One, the method of initial rates, involves measuring the rate at the beginning of the reaction as a function of the concentration of each reactant in turn. If, for example, it is found that the initial rate is twice as large when the starting concentration of reactant A is made twice as big, then it follows that the order of the reaction for A is 1 (that is, the exponent to which the [A] must be raised in the rate law is 1). This method suffers from a number of disadvantages, and is not widely applicable. However, we will find it useful in this experiment.
A second method, very commonly used, is the pseudo-order method. Applied to reaction (2), it would go something like this. There are only two reactants, A and B, and we assume that A has some physical property (such as light absorbance in the visible region of the spectrum) that enables its concentration to be monitored. One then conducts a kinetics study in which the initial concentration of A is appropriate to give a reasonable value of the property being measured, and the initial concentration of B is at least 10 times larger than that of A. It is thus guaranteed that over the course of reaction, during a time period in which the amount of A substantially decreases, the concentration of B changes hardly at all; i.e., [B] remains essentially constant over the course of the reaction. In this case the rate law in (3) takes a simplified form in which the rate depends only on [A]:
Regardless of the value of n, equation (4) is integrable to give the following results, where in all cases [A]o signifies the initial concentration of A:
| n | Integrated Rate Law |
|---|---|
| 1 | ln([A]/[A]o) = -kt |
| 2 | 1/[A] = 1/[A]o - kt |
| 3 | 1/[A]2 = 1/[A]o2 - kt |
The first 2 cases, n = 1 and 2, are by far the most common. When n=1, the reaction is said to be pseudo-first-order in A, and is pseudo second order when n=2. The prefix "pseudo" is meant to convey that there may also be a rate dependence on [B], but that this is invisible because [B] is large.
To determine the order in [B] one then repeats the study using a different, still large, concentration of B. The values of kobsd can then be plotted versus the concentrations of [B] to determine the dependence on [B].
In this experiment you will use the method of initial rates to determine the rate law for reaction (6):
Click here for structures of the reactants and product.
The reaction is catalyzed by hydroxide ion, OH-, and does not proceed at an appreciable rate unless hydroxide is added. When run in a mix of ethanol and water as solvent, the product begins to precipitate from solution shortly after the start of reaction. The precipitation manifests first as cloudiness in the solution that is readily visible. We will use the time elapsed between start of reaction and the first appearance of cloudiness as a measure of the initial reaction rate.
In all kinetics runs we will use a total reacting solution volume of 4 mL, prepared by mixing judiciously chosen volumes of solutions of benzaldehyde, acetone, and sodium hydroxide. In some cases, we will have to add extra ethanol or water to make the volume up to a total of 4 mL. Because reaction (6) runs smoothly only when the two reactants are present in the correct stoichiometric ratio, we will not be able to independently vary the initial concentrations of benzaldehyde and acetone. Rather, we will change both at the same time to maintain the stoichiometric ratio, and make an educated guess at the rate dependence on their concentrations.
Finally, we will qualitatively study the dependence of reaction rate on temperature.
Focus Questions
Equipment and Materials
Note to instructor: Click here for recipes for preparation of solutions. The more accurate the concentrations, the more reproducible and internally consistent the results will be. Please prepare solutions carefully and accurately.
Safety
Safety glasses must be worn at all times in the laboratory. Sodium hydroxide is corrosive. Avoid contact with skin and clothing. In case of skin contact, rinse copiously with water. Benzaldehyde is classified as a toxic irritant. Avoid contact with skin.
Experimental
Record all data in your notebook. ESTABLISH a clean work space. Obtain the necessary equipment and clean the glassware thoroughly using brushes and Alconox detergent. Rinse with distilled water and dry thoroughly. You may organize your work space as shown here if you desire, or you may establish your own organization.
Some Tips for Improving Your Reproducibility. It is difficult to get reliable kinetics data without a consistent, reproducible protocol (a systematic way of doing things). There are several pitfalls in this experiment that we will address here so that you can avoid them:
The Rate Law at Ambient Temperature. Measure room temperature and record it.
In separate clean, dry 50-mL Erlenmeyer flask, obtain 10 mL of 2.0 M benzaldehyde solution, 10 mL of 1.0 M acetone solution, and 10 mL of 3.0 M NaOH solution.
Into a 10- or 25-mL Erlenmeyer flask, pipet 1.00 mL of acetone solution and 2.00 mL of NaOH solution. Swirl to mix. Then pipet 1.00 mL of benzaldeyhyde solution, QUICKLY "inject" it into the acetone/NaOH mixture while starting the timer, IMMEDIATELY swirl the solution, and place the flask on a black background. Stop timing at the first appearance of white cloudiness in the solution. (NOTE: IF POSSIBLE, SWIRL WHILE QUICKLY ADDING THE BENZALDEYHYDE. YOU MAY SEE SOME INITIAL CLOUDINESS APPEAR WHERE THE NaOH ENTERS THE SOLUTION. HOWEVER, THIS SHOULD DISAPPEAR WITH SWIRLING. YOU ARE LOOKING FOR CLOUDINESS THAT APPEARS UNIFORMLY THROUGHOUT THE SOLUTION.) Record all concentration conditions and the reaction time. Repeat the experiment twice more, and average the times obtained. In all cases, pour the reaction mixture into a large beaker when finished. All subsequent reaction mixtures should also be poured into this beaker, which should be covered and kept in the hood.
Clean the reaction flask by flushing copiously with tap water. Rinse with distilled water, then use paper towelling to completely dry the inside of the flask.
Again in triplicate, carry out a rate experiment using 1.00 mL acetone solution, 1.00 mL of water, 1.00 mL of NaOH solution, and 1.00 mL benzaldehyde solution, transferred IN THAT ORDER. Again, it is essential to add the benzaldehyde solution QUICKLY, followed immediately by rapid swirling. Determine the average time to first appearance of cloudiness. Record all observations and data.
In triplicate, carry out a rate experiment using 0.50 mL of acetone solution, 1.00 mL of ethanol, 2.00 mL of NaOH solution, and 0.5 mL of benzaldehyde solution, transferred IN THAT ORDER. Determine the average time to first appearance of cloudiness. Record all observations and data.
The Rate Constant at Low Temperature. Make an ice bath and submerge your reactant solutions in it for 1-2 minutes. Record the temperature of the bath. Carry out the following experiment.
In a pre-cooled 10- or 25-mL Erlenmeyer flask, mix 1.00 mL of acetone solution and 2.00 mL of NaOH solution. Smoothly inject 1.00 mL of benzaldeyhde solution while starting the timer, rapidly swirl, and place in the ice bath. Watch closely for the first appearance of cloudiness. (NOTE: the cloudiness will probably develop more gradually than in the room temperature runs, so be alert.) Record the time.
The Rate Constant at Elevated Temperature. Make a 45 oC water bath and submerge your reactant solutions in it for 5 minutes. Record the temperature of the bath. Carry out the following experiment.
In a pre-warmed 10- or 25-mL Erlenmeyer flask, mix 0.50 mL of acetone solutions, 1.00 mL of ethanol, and 2.00 mL of NaOH solution. Then pipet in 0.5 mL of benzaldehyde solution while starting the timer, rapidly swirl, and place in the water bath. Watch closely for the first appearance of cloudiness.
The following procedures may be performed if time permits, at the discretion of the instructor.
Isolation of Reaction Product. The beaker containing the contents of all reaction solutions should at this point contain a yellow solid in a yellow-orange, cloudy liquid. Carefully pour off as much of the liquid as possible into the solvent waste bottle in the hood, leaving the solid behind in the beaker. In doing this, you may have to use your spatula to "hold back" the solid, similar to the way you pour water off just-cooked green beans.
Purification. Recrystallize your product from a small volume of acetone. First slurry the product with about 5 mL acetone. Add a little more acetone if necessary to dissolve all of the solid, but do not add a large excess. Suction filter the solution (or filter it through a Kimwipe-plugged Pasteur pipet) and transfer the filtrate to a small beaker. Add an equal volume of water to the beaker, then place the beaker in an ice bath until crystals form. Suction filter the mixture and suction dry.
Characterization of Product. Obtain the infrared spectrum of the product.
Clean-up. When you have finished all of your work:
Disposal Methods
Remaining benzaldehyde and acetone solutions should be poured into the organic waste bottle. Dibenzalacetone should be transferred to the large labelled bottle in the hood.
References
Preparation
Kinetics: The Reaction of Benzaldehyde and Acetone
