BROMINATION OF ACETONE

The form of the rate law for the bromination of acetone using an acid catalysts is determined. This experiment illustrates the fact that the rate law does not necessarily bear any simple relationship to the stoichiometric equation for the reaction

Apparatus

Spectronic 20 (L); set of 4 cells with caps (S); timer (S); nine 50 ml stoppered Erlenmeyers (preferably glass stoppered)(D,S); six 100 ml volumetrics (D,S); three 5 ml pipettes (D,S); 10 ml pipette (D).

Chemicals

Saturated bromine water (L) in hood; 100 ml 1 M HCl (L); 50 ml acetone (L); distilled H₂O (L)

_.

Use the Bausch and Lomb Spectronic 20 (or other suitable visible spectrometer) for making measurements. Determine the concentration of the stock bromine solution by measuring its absorbance at 450 nm and using the equation A = abc where A = measured absorbance, a = absorbency index of 100 M^-1cm^-1 at 450 nm for aqueous bromine solutions, b = cell thickness (1.0 cm), and c = concentration. Do this before preparing the stock solution of bromine H₂O. Check the concentration of the saturated bromine water with the spectrophotometer. If the instrument gives an off scale reading (as indicated by a flashing light) it is suggested that the saturated Br₂ solution be diluted by a factor of ten. Then the concentration should be measured again. If necessary repeat the process. From the on scale measurement the original concentration can be determined. Adjust the proportion accordingly to prepare the 0.03 M stock bromine water.

Now, prepare the following stock solutions:

1. 100 ml 6.0 M acetone in H₂O (44.05 ml acetone diluted to 100 ml with H₂O);

2. 100 ml 0.03 M bromine water (prepared as described above);

3. 100 ml 1.0 M HCl.

Take 20 ml (use pipette) of each stock solution and dilute with 20 ml H₂O to give 40 ml each of 3 M, 0.015 M, and 0.5 M acetone, bromine and HCl solutions, respectively.

Now, take 10 ml of each stock solution and dilute to 100 ml in volumetric flasks to give 100 ml each of 0.6 M, 0.003 M, and 0.1 M acetone, bromine, and HCl respectively.

CAUTION: Carry out the preparation of all these solutions in a HOOD! Keep all solutions stoppered throughout the experiment. One of the products of the reaction is a lachrymator. Poor laboratory technique will be most obvious, and will have your lab instructor in tears!

For all runs, use H₂O in the reference cell. Three sets of kinetic experiments are to be run, varying the concentrations of Br₂, acetone, and HCl independently.

Set one: 5 ml 6 M acetone + 5 ml 1 M HCl + 5 ml 0.003 M Br₂

0.015 M Br₂

0.030 M Br₂

Set two: 5 ml 0.6 M acetone

3.0 M acetone + 5 ml 1 M HCl + 5 ml 0.03 M Br₂

6.0 M acetone

Set three: 5 ml 6 M acetone + 5 ml 0.10 M HCl

0.50M HCl + 5 ml .03 M Br₂

1.00 M HCl

In each case mix the reactants in a 50 ml flask, adding the Br₂ solution last. Note the time of adding Br₂ as the starting time of the reaction. Mix well and transfer 3-4 ml to spectrophotometer cell to start measurements.

Since bromine absorbs strongly at the blue end of the visible spectrum, this reaction may be studied conveniently with a spectrophotometer such as that described before. The absorbency indices of Br₂ dissolved in distilled water are 160 M^-1 cm^-1 at 400 mm, 100 M^-1 cm^-1 at 450 mm, 30 M^-1 cm^-1 at 500 mm and

8 M^-1 cm^-1 at 550 mm. The solubility of Br₂ in water is about 0.21 at 25°, and it is more convenient to use water saturated with Br₂ in preparing solutions than to use pure liquid bromine. The concentration of the stock bromine solution is determined, using the absorbency indices given above. In accurate work it is necessary to determine cell corrections by intercomparing the various cells filled with solvent. In order to obtain the greatest percentage accuracy in the determination of Concentration, the absorbency should be in the range of about 0.2 to 0.7 (percent transmission of about 60 to 20). The spectrophotometer cells and all bromine solutions should be kept covered.

Hydrochloric acid is a convenient catalyst, and a suitable concentration range is 0.05 to 0.5 M. It is desired to determine the initial rate law, i.e., the order of the reaction with respect to acetone, hydrogen ion, and bromine, and the rate constant .

In order to avoid complications, only initial velocities are used. A series of experiments is designed to determine the concentration effects. The acetone concentration may be varied in the range 0.1 to 2 M. The bromine concentration may be varied in the range 0.001 to 0.01 M. Some of the higher concentrations should be run first because the initial velocities can be determined in a shorter time. Using a cell holder for four cells, three kinetic experiments can be run simultaneously. In order to obtain accurate results, the spectrophotometer should be equipped with a thermostating arrangement so that the reacting solutions can be held at the desired constant temperature. In the absence of this equipment, the temperatures of the solutions should be measured at the beginning and end of the experiment.

QUESTION: For each kinetic run, why is the Br₂ solution added last to the reaction mixture?

QUESTION: What is the method of initial rates?

THEORY.[1] The rate law for a reaction cannot be predicted from the balanced equation for the reaction. The rate law can only be determined experimentally. From the form of the rate law certain conclusions can be reached about the mechanism of the reaction.

The stoichiometric equation for the bromination of acetone is

O O

║ ║

CH₃CCH₃ + Br₂ Û CH₃CCH₂Br + Br^- + H⁺

The reaction rate increases with the concentration of H+ in acidic solution or with the concentration of OH- in basic solution. The balanced equation for the reaction occurring in acidic solution involves hydrogen ion as a product, and one would anticipate, therefore, an increasing reaction rate in the course of an experiment carried out in unbuffered solutions.

The rate of halogenation of acetone is independent of the concentration of

halogen, except at very high acidities.[2] The rates of reaction with the different halogens (chlorine, bromine, and iodine) are identical, and the same as the rate of racemization (for the case of optically active ketones), within a few percent.[3-5]. These facts can be accounted for in terms of the mechanism

Since ketones are very weak bases, the equilibrium in the first reaction is unfavorable for the formation of the ion reaction (I). Under these circumstances (I) = K(Acetone) (H⁺:B), where K is the equilibrium constant for reaction (1), :B is the acid's conjugate base, and the parentheses indicate concentrations [ ]. The rate equations for Enol and Product P are, according to the mechanism, the following

d(Enol)/dt = k₂(l) - [k_-2(H⁺:B) + k₃(Br₂)](Enol) (5)

d(P)/dt = k₃(Br₂)(Enol) (6)

These rate equations may be solved for d(P)/dt under steady state conditions by letting d(E)/dt = 0 and substituting (I) = K(Acetone)(H⁺). In this way we obtain

(7)

This equation takes on a simpler form if the Enol which is formed in the first two • steps reacts much more rapidly with halogen than with hydrogen ions; That is, k₃(Br₂) » k_-2(H⁺).

d(P)/dt =k₃K(Acetone)(H⁺)

This is in accord with the observation that the overall reaction is first order in ketone and acid but independent of the concentration of halogen. It can be seen that in general the apparent second-order rate constant is made up of a combination of k₁ the equilibrium constant K for the formation of ion I.

A great deal of research has been carried out on this particular reaction,

the above mechanism can be extended to include racemization, deuterium exchangers and catalysis by bases.[7] The overall reaction does not stop with the monobromo- acetone, but it is not necessary to consider the subsequent reactions if only initial reaction rates are studied.

Caution: Experimental operations involving bromine should be carried out in a well-ventilated hood, and solutions containing bromoacetone should be kept stoppered.

This reaction can also be studied titrimetrically. The reaction is stopped by

pipetting aliquots into sodium acetate solution to raise the pH. Potassium iodide is added, and the I₂ formed is titrated with sodium thiosulfate.

CALCULATIONS. The reaction rates in moles per liter of Br₂ reacting per second are calculated for each reaction mixture, and the approximate uncertainty is estimated. It is the initial velocity which is desired in each case. Plots of initial velocity versus (acetone), (H⁺), and (Br₂) are prepared to determine the order with respect to each of these substances. After the form of the rate law has been determined, the best value of the rate constant is calculated.

Suggestion for further work. The reaction of acetone with iodine may be studied. A 0.01 M KI solution is used as solvent because the solubility of iodine in pure water is so low.

The catalysis of the reaction by acetate-acetic acid buffers may be studied.[8,9] Since the acetate ion is also a catalyst, there will be terms in the rate law for both acetate ion and acetic acid. There is also a term proportional to the product of these concentrations. The reaction is subject to general acid-base catalysis.[l0]

The catalysis of the reaction by bases such as pyridine or hydroxyl ion may also be studied.[7]

References

1. L. P. Hammett, “Physical Organic Chemistry,” McGraw-Hill Book Company, New York, 1940.

2. L Zucker and L P. Hammett, J. Am. Chem. Soc., 61:2791 (1939).

3. Bartlett and C. H. Stauffer, J. Am. Chew. Soc., 57: 2580 (1935).

4. S. K. Hsi and C. L Wilson, J. Chem. Soc., 1936: 623.

5. C. K Ingold and C. L. Wilson, J. Chew., Soc., 1934: 773.

6. S. K Hsi, C. K Ingold, and C. L Wilson, J. Chem. Soc., 1938: 78.

7. P. D. Bartlett, J. Am. Chem. Soc., 56: 967 (1934).

8. R. P. Bell and P. Jones, J. Chem. Soc., 1953: 88.

9. H. M. Dawson and J. C. Spivey, J. Chem. Son, 1930: 2180.

10. H. M. Dawson, C. R. Haskins, and J. E. Smith, J. Chem. Soc., 1929: 1884.