Chemical Kinetics

Experiment 3

02-20-2016

Name: Tamer Almajid.

Section: 05, chem216

Introduction:

Chemical kinetics involves the study and discussion of chemical reactions with respect to reactions rates. The experiment was conducted to recognize certain factors affecting reaction rates and to determine and to explain their effects. The factors studied were concentration of reactant(s), temperature, and the presence of a catalyst. The experiment involved a variety of mixtures which were observed and their rates of  ‬reaction  ‬were compared ‬ and ‬ recorded. The reactant S2O8-2 was introduced into the reaction mixture in the form of (NH4)2S2O8, and I- was introduced into the reaction mixture in the form of KI. The balance net ionic equation is as follows:

1. S2O8-2 + 3I- → 2 SO4-2 + I3-

By using “clock” reaction: (2) 2S2O3-2 + I3- → S4O6-2 + 3I-

We could find the initial rate. The measurement involved is the time required for small portion of the reactants to be consumed.

The free triiodide ion, I3- , in reaction (1) is allowed to react with thiosulfate, S2O3-2 , no free triiodide can be there as long as thiosulfate is there in the solution. As soon as all the thiosulfate is gone, the I3- occurs. The presence of I3- is detected by its reaction by using starch indicator which is a deep blue color. The reactant S2O3-2 is introduced into the reaction mixture in the form of Na2S2O3

‬ It ‬  ‬was ‬ found ‬  ‬that  ‬acids,  ‬substances with  ‬higher concentration, higher temperature, more surface area exposed and the presence of a catalyst increases the rate of a chemical reaction. ‬‬‬‬‬‬‬‬‬‬‬‬‬

In this experiment every reaction mixture will have a fixed amount of S2O3-2 the starch indicator. As fast as the blue color appear, as fast as the reaction happen.

Discussion:

The partial reaction orders of this experiment was determined to be 1 for each reactant. This consistent with the data recorded in certain trials within the A and B experiments. In part A, from trials 3 to 2, the concentration of(NH4)2S2O8 is held constant while KI is doubled. Dividing the rates of KI from trials 2 by 3 along with their respective concentrations allows for the expression of 2 =2y. This was also the case for experiment B, trials 1 to 2. Once the partial reaction orders were determined for the rate law of this experiment, the rate constants for experiment A, trial 2 and experiment B, trial 3 was determined. Comparing these two values, the rate constants further support the partial reaction orders of each reactant. In experiment A, trial 2 the concentration of each reactant is held constant while in experiment B, trial 3, the concentration of (NH4)2S2O8 is reduced by half from 0.20 M to 0.10 M. When comparing the rate constants from A2 to B3, the constants nearly half as well from 3.3 x 10-5M/S to 1.8 x 10-5 M/S. In part C of the experiment, concentrations of the reactants were held constant while the temperature of the reaction was increased after each trial. As expected, the rate of reaction increased, resulting in an activation energy of 50.8 kJ/mol. The value of A was determined to be 1.9 x 10-4 M-1S-1. In experiment E, concentrations were held constant and the temperature was increased after each trial just like in experiment C, but this time a drop of 0.10 M copper nitrate was added into the reaction, resulting in generally quicker reaction rates in seconds and M/S. The activation energy of Experiment E was calculated to be 46.5 kJ/mol, just as a catalyst would expect to do, successfully lowering the activation energy required for the reaction. The value of A was then determined to be 2.2 x 10-5 M-1 S-1.