CRYSTAL VIOLET REACTION

 

Purpose

 

            The purpose of this experiment is to determine how the rate constant for a reaction changes with ionic strength of the solution and with temperature.

 

INTRODUCTION:  The experiment follows a procedure described in J. Chem. Ed., 41, 48 (1964).  By using the UV-VIS spectrophotometer the rate constant k is measured for the reaction

 

 

Here  is crystal violet, a violet-colored organic dye of molecular weight 408.0, which reacts with OH- to form a colorless complex.  The reaction is studied at three temperatures to determine the activation energy, and at three ionic strengths (at one temperature) to determine the effect of inert electrolytes on the rate constant.

 

THEORY:  We assume the rate law for reaction (1) is

 

 

If , then  remains approximately constant during the reaction, and

 

 

where

 

Integration of (3) between  and   yields

 

If concentrations are determined by a spectrophotometer we can use the relation

 

 

where A is the absorbance.

 

Using  (5)  and  (6)  we obtain

 

 

Thus  can be determined from a plot of  ln  vs. t; because (OH-) is known, the rate constant k can then be calculated from (4).

 

            According to transition-state theory the rate constant at a given temperature for a bimolecular reaction varies with the total ionic strength, m,  of the system,

 

 

mcan be calculated from the charges, , and the molarities, , of all ions in the system:

 

 

The value for  B  can be determined (if the plot of   vs.  is linear) and compared with the theoretical value, thereby testing the validity of the theory for this system.  (B is the slope.)

 

The Arrhenius activation energy, , appears in the equation  

 

 

From a plot of  vs.  (absolute temperature), at constant ionic strength,   can be determined if the plot is linear.  The slope is .

 

 

EQUIPMENT AND CHEMICALS

 

            Visible spectrophotometer (P. E. Lamba 3, Turner 350, Coleman 124)

            Crystal violet (0.03 g/l),

            Electrolyte solutions may be prepared by dissolving a quantity of KNO3 in

            0.008M NaOH (For example 0.002 mole KNO3 dissolved in 50 ml 0.008M

            NaOH.)

 

PROCEDURE:  A stock solution containing 0.030g of CV+.Cl- per liter is available.  Part A of the experiment requires the determination of k at three different ionic strengths but at one temperature, 25°C.  For each determination, 50 ml. of 0.006g./l solution of CV+Cl- is mixed briefly with 50 ml. of an electrolyte solution; a portion of the mixture is then placed in the sample cell ( water is in the other cell), and values for absorbance of solution are recorded every minute for about 15 minutes.The spectrophotometer wavelength is set at the position of highest absorbance (about 586 nm) as determined with a non reacting solution of CV+Cl-.  The first 50 ml. portion of electrolyte used is 0.008 M in NaOH; the portion used for the second run is 0.008 M in NaOH and 0.04 M in KNO3; for the third run the concentrations in the 50ml. portion are 0.008 M NaOH and 0.16 M KNO3.  (Note: You cannot get a solution that is 0.008M NaOH and 0.04M KNO3 by mixing 0.008M NaOH and 0.04M KNO3. Since the solutions are diluted 1:1 in each other, concentrations are reduced by one-half.)

 

            In part B of the experiment, k is determined at the temperatures 25°C, 30°C, and 35°C, with the 50ml. electrolyte portion being 0.008 M in NaOH and no KNO3 present.  As before the other 50 ml. portion to be mixed contains 0.006g. of CV+Cl-/l.

 

 

CALCULATIONS:  For each run calculate and  at every time t.  Plot  vs. t and determine ; then calculate k from equation 4.  For  runs in part A calculate and determine B in equation 9.  Remember, the 50 ml. electrolyte portions were diluted 1:1.  From the runs in part B determine  in (11).  Prepare tables and graphs displaying the above data and results.  Discuss the validity of equations 8 and 11 for this system.