Kinetics in chemistry deals with the rate at which a chemical reaction occurs. This rate, which is referred to as the reaction rate, is defined as the change in concentration of a reactant or product with time, and is measured in M/s. The rate of a reaction is proportional to the concentration of reactants. An equation called the rate law expresses the relationship of the reaction rate to the rate constant, k, and the concentrations of the reactants raised to some powers, x and y, found experimentally. The rate law is expressed as, rate = k [A]x[B]y. The constant k is equal to the rate divided by the concentration of a certain substance. The purpose in this lab was to experimentally determine the rate constant k, as well as the exponential values of x and y in the rate law.
The procedure of this lab was obtained from the student laboratory course website.
Table 1: Reagent Volumes
|1||0.5 mL||0.8 mL||0.8 mL||1.9 mL||4.0 mL|
|2||0.5 mL||1.6 mL||0.8 mL||1.1 mL||4.0 mL|
|3||0.5 mL||0.8 mL||1.6 mL||1.1 mL||4.0 mL|
Table 2: Initial Concentrations, Times, and Rate for Each Trial
|Trial||[I2]||[Acetone]||[H+]||Time (sec)||Rate [I2]||k|
|1||6.25e-2 M||0.68 M||0.2 M||184 sec||3.40e-6 M/sec||2.50e-5 M-1s-1|
|2||6.25e-2 M||1.36 M||0.2 M||125 sec||5.00e-6 M/sec||1.84e-5 M-1s-1|
|3||6.25e-2 M||0.68 M||0.4 M||88 sec||7.10e-6 M/sec||2.62e-5 M-1s-1|
To conduct this experiment, the groups placed 1.9 mL of distilled water, 0.8 mL H+ (HCl), and 0.8 mL of an acetone solution into a 4.0 mL cuvette. 0.5 mL of I2 was then added to the cuvette, and the group immediately started timing the reaction. They mixed the contents of the solution by inverting the cuvette several times before placing it into a calibrated spectrometer. The absorbance rate was monitored at 400 nm until it reached a nominal zero value. The time was then stopped, and the groups were able to determine the rate. Two more trials were conducted, first doubling the original volume of the acetone and keeping the H+ at a constant, then in trial 3, doubling the volume of the H+ and using the original volume of acetone – 0.8 mL. Constantly keeping the volume of I2 at 0.5 mL, and doubling one solution while keeping the other constant made it possible to later calculate the value of the rate constant, k. The chemical reaction being studied was chemical kinetics—the rate at which I2 disappeared. To determine the rate of disappearance of I2 in the reaction, the equation M1V1=M2V2was used to find the concentration of I2. Then, that value was divided by the time elapsed to result in the rate. When I2 was first added to the cuvette, it was dark red in color. As the reaction progressed, the solution lost its color and became clear, consuming the I2 completely. At this point, the spectrophotometer displayed that at 400 nm, zero light was being absorbed in the solution. The starting concentrations were varied according to the experiment design in order to calculate the rate law exponents. The rate law does not include [I2] because I2 does not impact the rate of chemical reaction under the conditions selected. It isn’t included because the rate of disappearance of I2 was what was being solved for. The rate law determined from this lab is rate= k[Acetone][H+]. By using the values of [Acetone], [H+], and the rate, and taking the average of the three trials, the value of k was found to be 2.32e-5 M-1s-1.
The goal of determining the values of k, the exponents x and y, and the rate of disappearance of I2 were successfully met. The expected results matched up with the obtained results—the concentration of acetone and H+ are directly related to the rate of reaction. Both [Acetone] and [H+] are first order reactions, resulting in an overall second order reaction. The main source of error in lab came from not correctly measuring out the substances, resulting in a very askew time and rate of reaction. Having incorrect amounts of each solution in the cuvette directly affected the rate at which I2 disappeared, which in turn made the results not as clear or concise.
The goal of this lab was to experimentally determine how the concentrations of Acetone and H+ affect the rate at which I2 disappears in a reaction by calculating the values of k, the exponents x and y, and putting those values into the rate law equation. Those numbers were found by altering the amounts of either acetone or H+ used in each trial, which made it possible to use an equation provided above to solve for the value of k, and the exponents, x and y, in the rate law equation. It was found that both Acetone and H+ have a direct effect on the reaction rate of I2. The rate law for acetone iodination is rate= k[Acetone][H+]. The average value of k calculated from the three trials was found to be about 2.32e-5 M-1s-1.
2 Chemical reactions can occur in many different rates of time. Temperature and concentration affect the rate properties/chemical kinetics of the reactions. The purpose of the experiment is to evaluate the rate constant k and the reaction orders m, n, and p . Once these values are determined then a predicted rate and time can be determined. The rates depend on the concentrations of the Acetone, H + ,and I 2 ions according to the rate law. Introduction The concentration of the reactants can influence the rate in which the reaction takes place. The amount of concentration and the rate of reaction have a relationship where high concentration will make a high reaction time, or vice versa. This experiment studies rate properties of the reaction between iodide ion and acetone ion under acidic conditions. The rate law can be used to determine the main purpose of the experiment, which is to evaluate the rate constant k and the reaction orders m, n, and p using the following equation. (1) Rate= k (Acetone) M (H + ) n (I 2 ) p In the above formula, k is the rate constant. Acetone, H + ,and I 2 are the concentrations, which is the molarity of the reactants in the experiment. The values of m, n, and p are either 0, 1 st , or 2 nd order. In the zero order the concentration has no effect on the solution. If it is in the first order then the concentration is directly proportional. In the second order the change in the rate of the concentration is effected exponentially and in the third order even a very slight effect has an exponential, exponential effect on the rate of concentration. In order to get the concentrations of each solution in total volume ( in this experiment it is 50mL) we must take the molarity of each individual solution, multiply it by amount of that solution in mL, and then divide by total amount of all solutions in flask (50 mL). Example: (Molarity of one solution X mL of solution)/ (Total amount of all solutions in mL) = solution concentration