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The Rate Law For Chemical Reaction Among Hydrogen Peroxide, Iodide, And Acid

3073 words - 12 pages

The Rate Law for Chemical Reaction Among Hydrogen Peroxide, Iodide, and Acid

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

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...

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