Introduction:
The goal of this experiment is to determine the rate of a reaction by way of kinetics. This is accomplished in a reaction between Crystal Violet and Sodium Hydroxide by measuring the concentration of the reaction over time, via a spectrophotometer. Then, by making use of the rate law, the concentrations of the chemicals in the reaction, and the change in time, we can calculate the rate. This equation is the rate law:
Then, as we are making use of absorbance, we must also make use of the Beer–Lambert law, to solve for the concentration:
A=εlc
(A=Absorbance, ε=the molar absorption coefficient, l=the path length through the solution that the light must travel, c=concentration)
Thus:
c=A/(εl)
Then, we must calculate the order of the reaction, as the equations used differ depending on the order of the reaction, as depicted in the table below:
Characteristics of Reaction Order:
One way to go about finding the order to which the reaction applies is to then measure the initial rate of reaction under varying concentrations of the reactants.
Which then becomes:
Then, if we set it so that the second experiment (differentiated with the subscript 2, as compared to experiment 1) uses double the initial concentration of one of the reactants (in this case OH-), we can simplify to the following equation, and solve for n:
Then, by repeating the process for the Crystal Violet in a third experiment, we can solve for m:
Finally, in knowing this, we can calculate the rate, and, then, with that, using the equation for the particular order, the rate constant.
Equipment:
- Data Collection System
- Spectrophotometer (Red Tide)
- Sensor Extension Cable
- Cuvette
- Beaker (3), 50-mL
- Syringe (3), 1-mL
- Plastic Covers
- 0.1 M Sodium Hydroxide, 20 mL
- 1.2e-5 M Crystal Violet, 20 mL
- Water, Deionzed, 30 mL
- Marking Pen
- Kimwipes
Methods:
- Gathered all equipment and materials. Inspected all equipment for damage or contamination.
- Set up Data Collection System.
- Calibrated the Spectrophotometer.
- Put 20 mL of Crystal Violet solution in beaker. Repeated with the distilled water, and the Sodium Hydroxide. Covered Sodium Hydroxide with the plastic covering.
- Labeled each syringe after each solution. Filled according to Table 1, taking care to remove any air bubbles.
- Table 1: Initial concentrations and Volumes (Note: "E" is simply for scientific notation)
- Inserted specified quantity of water and NaOH into clean cuvette, before beginning the data recording process, injecting the Crystal Violet solution, and inverting it rapidly.
- Immediately wiped the cuvette with a Kimwipe, then inserted it into the Red Tide Spectrophotometer.
- Recorded data until absorbance had decreased below 0.1.
- Saved data, then repeated process two more times, following the concentrations and volumes from Table 1 for trials 2 and 3 respectively.
- Disposed of solutions in designated containers. Cleaned all glassware and cuvettes, set them to dry in their designated areas, and returned all equipment to its point of origin.
Data Analysis:
Graph 1: Concentration over Time (M/s)
From Left to Right, slopes for 1, 2, and 3.
Initial Rates (M/s)
Trial 1: 1.52E-08
Trial 2: 4.37E-08
Trial 3: 2.44E-08
n=0.841
m=0.683
Thusly, pseudo-first order.
Then, using the rate formula for pseudo-first order equations:
Table 2: Determination of the reaction rate constant:
The average rate constant (in inverse seconds) is:
0.207 (1/s)
Sample Calculation:
Initial Rates (M/s)
Trial 1: 1.52E-08
Trial 2: 4.37E-08
Trial 3: 2.44E-08
Calculations of m and n (rates are in ((1/M)(1/s))):
Determination of the reaction rate constant:
Trial 1: ln(0.019)=ln(0.051)-k192.5
Trial 2: ln(0.043)=ln(0.102)-k83.5
Trial 3: ln(0.026)=ln(0.081)-k216.5
Calculating for the rate constant (k' is in (M/s), Concentration of OH- is in M.):
The average rate constant (in inverse seconds) is:
0.207 (1/s)
Conclusion:
The goal of this experiment is to determine the rate of a reaction by way of kinetics. This is accomplished in a reaction between Crystal Violet and Sodium Hydroxide by measuring the concentration of the reaction over time, via a spectrophotometer. Then, by making use of the rate law, the concentrations of the chemicals in the reaction, and the change in time, we can calculate the rate. This equation is the rate law:
Then, as we are making use of absorbance, we must also make use of the Beer–Lambert law, to solve for the concentration:
A=εlc
(A=Absorbance, ε=the molar absorption coefficient, l=the path length through the solution that the light must travel, c=concentration)
Thus:
c=A/(εl)
Then, we must calculate the order of the reaction, as the equations used differ depending on the order of the reaction, as depicted in the table below:
Characteristics of Reaction Order:
One way to go about finding the order to which the reaction applies is to then measure the initial rate of reaction under varying concentrations of the reactants.
Which then becomes:
Then, if we set it so that the second experiment (differentiated with the subscript 2, as compared to experiment 1) uses double the initial concentration of one of the reactants (in this case OH-), we can simplify to the following equation, and solve for n:
Then, by repeating the process for the Crystal Violet in a third experiment, we can solve for m:
Finally, in knowing this, we can calculate the rate, and, then, with that, using the equation for the particular order, the rate constant.
The average rate constant (in inverse seconds) is:
0.207 (1/s)
(No calculation of error possible, as accepted rate of reaction was not given)
Possible sources of error include:
- Failure to properly clean the cuvettes.
- Failure to properly mix (via shaking) the water, NaOH and Crystal Violet solution in the cuvettes, prior to taking measurements.
- Contamination of cuvettes while handling them (smudging the exterior, etc).
- Contamination of solutions used, thus altering them from their formula and concentration.
- Failure to properly handle the solutions, introducing the likelihood of contamination (NaOH can react with carbon dioxide in the air), thus disrupting measurements.
- Contamination or other error, relating to the spectrometer (such as smudging the lens).
- Failure to generate a proper graph, due to lack of knowledge of the directions.
- Human error is always in effect, given that the laboratory does not function under ideal conditions. As such, there is always the possibility of inaccuracies with measurement, perception of measurement, inaccuracies of equipment, and other such errors. (However, this is not likely to be the sole cause of the inaccuracies within this experiment, though it may contribute to it.)
Possible improvements to this experiment include testing each cuvette immediately upon combining the solutions, using crystal cuvettes, slowing down the reaction (to get a smoother curve), using a more accurate spectrometer, wearing gloves while handling the cuvettes, keeping the NaOH properly covered, ensuring that lab partners read the experiment beforehand (increasing familiarity with the procedure, and minimizing human error), and repeating the experiment multiple times (to average out their absorption spectra, and minimize the impact of an anomalous result).