**Introduction:**

**Iodination of Acetone Lab Report**

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 iodination of acetone 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 of this lab was to experimentally determine the rate constant k, as well as the exponential values of x and y in the iodination of acetone rate law.

**Experimental:**

The procedure of this lab was obtained from the student laboratory course website. The iodination of acetone lab report involved preparing various solutions and monitoring the reaction rates of the iodination of acetone mechanism.

**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 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 the kinetics of the iodination of acetone—the rate at which I2 disappeared. To determine the rate of disappearance of I2 in the reaction, the equation M1V1=M2V2 was 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 a 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.

## Errors and Improvements

One major factor that affected the result of this experiment was the strength of the sodium hydroxide and sodium solutions. If either of these substances is left open in the atmosphere, they begin to lose their strength. During the experiment, the sodium hydroxide and sodium were both left open to interact with the environment for some time. Thus, the final answer did not match the theoretical value accurately because the strength was weakened, meaning that the numbers used to calculate the molar concentration were not as accurate.

In addition, the equipment used could have also contributed to the error as all pieces of apparatus have an uncertainty attached to them. These uncertainties are then applied to calculations in order to keep up the amount of uncertainty associated with the amount of material used. This uncertainty was found to be ±3.87% for all of the experiments. This value includes both uncertainties regarding those that applied when the solutions were made (e.g., the uncertainty of the mass balance that was used to measure the amount of sodium that was needed to make the sodium hydroxide) and the transfer of the solution from one instrument to another.

Human judgment also accounts for some of the error in this experiment as the person performing the experiment was required to read off many measurements from the pipette and burette. This error can be reduced by always ensuring that readings are always made at eye level and that the same person taking the readings is constant as judgment varies with each person.

## Conclusion

In conclusion, this experiment found the molar concentration of acetic acid in vinegar to be 0.44mol/L. However, this value was 50% inaccurate due to errors that occurred while conducting the investigations. The kinetics of the iodination of acetone were thoroughly explored, and the iodination of acetone rate law was determined to be 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.

By examining the rate law and the iodination of acetone mechanism, this lab report successfully highlighted the complexities and precise nature of chemical kinetics. Despite the experimental errors, the overall findings were consistent with the theoretical expectations of the iodination of acetone and its related kinetics.

Hi, do you know what is the theoretical rate constant for this reaction?

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