To determine the rate law for a chemical reaction among hydrogen peroxide, iodide and acid, specifically by observing how changing each of the concentrations

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Experiment 3 – Chemical Kinetics

Objectives

1.        To determine the rate law for a chemical reaction among hydrogen peroxide, iodide and acid, specifically by observing how changing each of the concentrations of H2O2, and H+ affects the rate of reaction.

2.        To observe the effects of temperature and catalyst on the rate of reaction.

Introduction

        Generally, two important questions may be asked about a chemical reaction:

(1)How far do the reactants interact to yield products, and (2) how fast is the reaction? “How far?” is a question of chemical equilibrium which is the realm of chemical thermodynamics. “How fast?” is the realm of chemical kinetics, the subject of this experiment.

In this experiment we will study the rate of oxidation of iodide ion by hydrogen peroxide which proceeds according to the following reaction:

H2O2 (aq) + 2 I-(aq) + 2H+(aq)I2(aq) + 2H2O(l)

By varying the concentrations of each of the three reactants (H2O2, I- and H+), we will be able to determine the order of the reaction with respect to each reactant and the rate law of the reaction, which is of the form:

Rate = k [H2O2]x[I-]y[H+]z

By knowing the reaction times (t) and the concentrations of H2O2 of two separate reaction mixtures (mixtures A & B), the reaction order of H2O2, x, can be calculated.

x = log(∆t2/ ∆t1) / log ( [H2O2]1/[H2O2]2 )

The same method is used to obtain the reaction order with respect to I- (mixtures A & C) and H+ (mixtures A & D).

Procedures

Part I) Standardization of H2O2 Solution

1.        A stand, a burette clamp and a white tile were collected to construct a titration set-up.

2.        A burette was rinsed with deionized water and then with 0.05 M Na2S2O3 solution.

3.        The stopcock of the burette was closed and the sodium thiosulphate solution was pour into it until the liquid level was near the zero mark. The stopcock of the burette was opened to allow the titrant to fill up the tip and then the liquid level was adjusted near zero.

4.        The initial burette reading was recorded in Table 1.

5.        1.00 cm3 of the ~0.8 M H2O2 solution was pipetted into a clean 125 cm3 conical flask.

6.        25 cm3 of deionized water was measured with a 50 cm3 measuring cylinder. It was pour into the conical flask.

7.        10 cm3 of 2.0 M sulphuric acid was measured with a 10 cm3 clean measuring cylinder. It was pour into the conical flask.

8.        1 g of solid KI (record the exact mass) and 3 drops of ammonium molybdate catalyst were added into the conical flask.

9.        The solution mixture was stirred until the KI dissolves.

10.        The reaction mixture was titrated in the conical flask with the sodium thiosulphate solution until it just turns pale yellow.

11.        3 drops of freshly prepared starch solution were added to the conical flask.

12.        The titration was continued until it just changes from dark blue to colorless. The final reading was recorded in Table 1. It was a first trial titration to estimate the volume of the sodium thiosulphate solution required. The volume of the sodium thiosulphate solution added in titration was calculated.

13.        The given sodium thiosulphate solution was added to the burette through a filter funnel if the volume remained was not enough to carry out another titration.

  1. Steps 6-13 were repeated to obtain 2 sets of consistent results. However, sodium thiosulphate solution was stopped draining at about 3 cm3 less than the estimated value. Then the sodium thiosulphate solution was added drop by drop until the reaction mixture in conical flask just changes from dark blue to colorless.

Part II) Reaction Rate Measurements

Six reaction mixtures will provide the information necessary to determine the effects of the concentrations of H2O2, I- and H+ on the rate of the reaction, as well as the effect of temperature and a catalyst.

Table 2 specifies the temperature and reagent volumes to be used for each reaction mixture.

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1.        All the reactants were placed in mixture A in the order in which they appeared in Table 2 except the H2O2 in a 250 cm3 conical flask.

2.        The temperature of the solution mixture was read and recorded.

3.        Specific amount of standardized H2O2 solution (according to mixture 1 in Table 2) was pipetted to the conical flask and stir continuously.

4.        The stop-watch was started when half of the H2O2 solution had been drained from the pipette.

5.        The conical flask was swirled to mix the contents well.

6.        Carefully watched the solution mixture for the sudden appearance of the blue color, ...

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