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 t = t and t = 0 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 k' can be determined from a plot of ln A_t 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,

m can be calculated from the charges, z_i, and the molarities, c_i, of all ions in the system:

The value for B can be determined (if the plot of ln k 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, E_a, appears in the equation

From a plot of ln k vs. (absolute temperature), at constant ionic strength, E_a 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^oC. 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^oC, 30^oC, and 35^oC, 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 A_t and at every time t. Plot ln A_t vs. t and determine k'; 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 E_a in (11). Prepare tables and graphs displaying the above data and results. Discuss the validity of equations 8 and 11 for this system.