**Introduction:**

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.

**Experimental:**

The procedure of this lab was obtained from the student laboratory course website.

**Results:**

Table 1: Reagent Volumes

Trial | I_{2} | Acetone | H^{+} | dH_{2}O | Total Volume |

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 | [I_{2}] | [Acetone] | [H^{+}] | Time (sec) | Rate [I_{2}] | k |

1 | 6.25e-2 M | 0.68 M | 0.2 M | 184 sec | 3.40e-6 M/sec | 2.50e-5 M^{-1}s^{-1} |

2 | 6.25e-2 M | 1.36 M | 0.2 M | 125 sec | 5.00e-6 M/sec | 1.84e-5 M^{-1}s^{-1} |

3 | 6.25e-2 M | 0.68 M | 0.4 M | 88 sec | 7.10e-6 M/sec | 2.62e-5 M^{-1}s^{-1} |

__Sample Calculations: __

**Discussion: **

** **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 I_{2 }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 I_{2} 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 I_{2} disappeared. To determine the rate of disappearance of I_{2} in the reaction, the equation *M _{1}V_{1}=M_{2}V_{2} *was used to find the concentration of I

_{2}. Then, that value was divided by the time elapsed to result in the rate. When I

_{2}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 I

_{2 }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 [I

_{2}] because I

_{2}does not impact the rate of chemical reaction under the conditions selected. It isn’t included because the rate of disappearance of I

_{2}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

^{-1}s

^{-1}.

The goal of determining the values of *k, *the exponents x and y, and the rate of disappearance of I_{2 }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 I_{2} disappeared, which in turn made the results not as clear or concise.

**Conclusion:**

** **The goal of this lab was to experimentally determine how the concentrations of Acetone and H^{+} affect the rate at which I_{2} 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 I_{2}. 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^{-1}s^{-1}.

## Leave a Reply

2 Comments on "Kinetics Lab Explained: Iodination of Acetone"

Hi I’m not too sure how you worked out the initial concentration of the iodine could you please explain?

The equation M1V1=M2V2 was used.

M1 = Concentration of the stock solutions

V1 = Volume of stock solution used

V2 = Total volume

M2 = Initial concentration

For example, to calculate the initial concentration, M2 of I3- for mixture I, we insert the value into the equation.

M1V1=M2V2

(0.02L) (0.023M) = (0.05M) M2 ∴ M2 = 9.2 x 10-3 M