Kinetics: The Reaction of Sn^II with Methyl Orange

• To explore the kinetics of a slow redox reaction
• To carry out a kinetics study under pseudo-order conditions
• To determine the rate law for the reaction

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

• The chemical nature of reactants and products
• The concentrations of reactants and products
• The temperature
• The presence of and concentration of a catalyst

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 numbers 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. It is utilized in the experiment, "Kinetics: The Reaction of Benzaldeyhde and Acetone."

A second method, very commonly used, is the pseudo-order method. Applied to reaction (1), 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:

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 word "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 k_obsd can then be plotted versus the concentrations of [B] to determine the dependence on [B].

In this experiment you will use the pseudo-order method to determine the rate law for reaction (6):

This is an example of an oxidation-reduction (redox) reaction, which are reactions in which electrons are transferred from one species to another. The indicator, methyl orange, has a strong absorbance in the visible region of the spectrum that we will use to monitor its disappearance. From the manner in which the amount of reactant decreases with time, the pseudo-order rate dependence on methyl orange will be determined. By running the experiment at a number of Sn^II concentrations, the rate dependence on [Sn^II] can be determined.

1. Which reactant undergoes a change in concentration during reaction?
2. Which reactant maintains constant concentration throughout the reaction?
3. The reaction is run under pseudo-order conditions in which reactant?
4. What is the order of reaction with respect to methyl orange?
5. What is the order of reaction with respect to tin(II)?
6. What is the value of the rate constant for the reaction at room temperature (including units)?
7. Describe in detail the experiments you would do in order to determine whether reaction rate depends on pH.

• 3 2-mL graduated pipets
• Syringe pipet pump
• 3 25-mL Erlenmeyer flasks
• 3 Pasteur pipets
• Pasteur pipet bulb
• Plastic spectrometer cell
• electronic absorption spectrometer
• Stock Solution of methyl orange indicator
• Stock Solution of SnCl₂.2H₂O

Safety glasses must be worn at all times in the laboratory. You will be working with 1 M acid solutions, which will burn the skin and clothes. In case of skin contact, flush the affected area with copious amounts of running water. Tin salts are toxic. Avoid ingestion and skin contact.

Record all data in your notebook. Obtain the necessary equipment and clean the glassware (but NOT the spectrometer cell) thoroughly using brushes and Alconox detergent. Rinse with distilled water and dry thoroughly.

Preparation of Solutions

Develop a procedure for the preparation of 10 mL of 0.001 M methyl orange indicator in water containing 1.0 M HCl. Have your procedure checked, then carry it out. A stock solution containing 0.20 M SnCl₂.2H₂O is available in the lab. Please obtain no more than 5 mL of this stock solution. Tightly stopper the solution containers when not in use to minimize contact with oxygen from the air.

Kinetics.

You will perform at least four kinetics runs, each with a different concentration of Sn^II. In the course of your work, you will record a number of absorbance-wavelength and absorbance-time spectral scans. Please save each one with a separate file name so that you can recall all of your scans at a later time. Before beginning your kinetics studies, you must determine the wavelength of light absorbed by the methyl orange reactant. Transfer 0.5 mL of your methyl orange solution and 1.5 mL of 1.2 M HCl to the cell, and record an absorbance-wavelength scan. Based on this scan, choose a wavelength at which to monitor the kinetics.

Please use the following procedure for each kinetics run. Refer to table 1 for amounts of water and tin solution to add to the cell.

1. Set up the spectrometer for an absorbance-time scan at your chosen wavelength. A scan time in the range of 300 to 1000 s should be used. You may want to do a trial run visually to determine the best scan time, which should be chosen as the time required for the methyl orange color to fade completely.
2. Using a 2-mL graduated pipet and a syringe pipet pump, transfer 0.50 mL of methyl orange solution to the spectrometer cell.
3. Again using a 2-mL graduated pipet and a syringe pipet pump, transfer the specified amount of 1.2 M HCl for run 1 (Table 1).
4. Transfer the run-1 amount of Sn^II solution to the cell, then QUICKLY stopper and shake the cell, place it in the electronic absorption spectrometer, and begin the absorbance-time scan.
5. Obtain the final absorbance of the solution. The final absorbance is the absorbance when reaction is finished.

When you have finished all runs, recall the absorbance-time files, overlay them, and print the overlay. On a separate sheet, print the spectrum of methyl orange in 1.2 M acid. Export all recorded spectra to a disk so that you can import them to a spreadsheet. (See the instructor for assistance with this if you need it.)

Clean-up. When you have finished all of your work:

• Clean all glassware with Alconox and water and shake dry.
• Rinse Pasteur pipets by inserting the nozzle of your wash bottle in the top end and squirting water through. Aspirate dry.
• Rinse the spectrometer cell thoroughly with distilled water, aspirate dry, and return to the labkit.
• Return all components to the labkit.
• Return all borrowed equipment to the instructor before leaving lab.

All solutions containing tin should be poured into the heavy metal waste jar in the hood. Solutions containing only methyl orange and acid (no tin) should be poured into the beaker provided in the hood.

References

1. GP Haight, J. Chem. Ed 1965, 42, 478.

1. This experiment.
2. The appropriate sections of your textbook.

1. The following data are obtained for the reaction below. 2A + 3B ---> 2D + F

   time, minutes	[A], molar	[B], molar
   0	1.6 x 10-3	0.5
   10	0.8 x 10-3	0.5
   20	0.4 x 10-3	0.5
   30	0.2 x 10-3	0.5

   In a second run, the following data were obtained

   time, minutes	[A], molar	[B], molar
   0	1.6 x 10-3	1.0
   2.5	0.8 x 10-3	1.0
   5	0.4 x 10-3	1.0
   7.5	0.2 x 10-3	1.0

   What is the rate law for the reaction?
   What is the value of the rate constant?
   What are the units of the rate constant?