Illustrated Guide to Home Chemistry Experiments

Chapter 12 Laboratory: Chemical Kinetics 223

LABORATORY 12.4:

dETERmINE THE EffECT of A CATALyST oN REACTIoN RATE

A catalyst is a substance that increases the rate

of a chemical reaction, but is not consumed

or changed by the reaction. A catalyst works

by reducing the activation energy needed to

initiate and sustain the reaction. For example,

two molecules of hydrogen peroxide can

react to form two molecules of water and

one molecule of molecular oxygen gas by the

following reaction:

2 H 2 o 2 (aq) → 2 H 2 o(l) + o 2 (g)

RIREEqU d EqUIpmENT ANd SUppLIES

£ goggles, gloves, and protective clothing

£ beaker, 150 mL (6)

£ graduated cylinder, 10 mL

£ graduated cylinder, 100 mL

£ pipette, mohr or serological, 10.0 mL

£ pipette, Beral (or other disposable)

£ burette, 50 mL

£ funnel (for filling burette)

£ ring stand

£ burette clamp

£ stopwatch

£ catalase enzyme solution (see Substitutions and

modifications)

£ hydrogen peroxide, 3% (50 mL)

£ potassium permanganate solution, 0.1 m (~150 mL)

£ sulfuric acid, 0.1 m (~250 mL)

£ distilled or deionized water

At room temperature, this reaction occurs very slowly because
few of the collisions between hydrogen peroxide molecules
have sufficient energy to activate the reaction. Furthermore,
commercial hydrogen peroxide solutions, such as the 3%
hydrogen peroxide solution sold in drugstores and the 6%
solution sold by beautician supply stores, are treated with
stabilizers (sometimes called negative catalysts) that increase
the activation energy for the reaction, further inhibiting it
from occurring.

If you add a catalyst to a solution of hydrogen peroxide, the effect
is immediately evident. The solution begins bubbling, as oxygen
gas is evolved. Numerous substances can catalyze the reaction of
hydrogen peroxide to water and oxygen gas, including many metal
oxides such as manganese dioxide, but the efficiency of catalysts
varies. One of the most efficient catalysts for hydrogen peroxide
is the enzyme catalase that is contained in blood. (Catalase
functions in the body as a peroxide scavenger, destroying
peroxide molecules that would otherwise damage cells.)

One catalase molecule can catalyze the reaction of millions of
hydrogen peroxide molecules per second. Immediately after each
pair of hydrogen peroxide molecules reacts, catalyzed by the
catalase molecule, that catalase molecule is released unchanged
and becomes available to catalyze the reaction of another pair of
hydrogen peroxide molecules. When all of the hydrogen peroxide
has reacted to form water and oxygen gas, you end up with as
many catalase molecules remaining as you started with.

In this lab session, we’ll measure the reaction rate of the
catalyzed reaction of hydrogen peroxide by adding a fixed amount
of catalase enzyme to measured samples of hydrogen peroxide.
After allowing the reaction to continue for measured periods of
time, we’ll stop the reaction by adding sulfuric acid to denature
(deactivate) the catalase and then titrate the resulting solutions

with a dilute solution of potassium permanganate to determine

how much unreacted hydrogen peroxide remains in each sample.

In acidic solution, the intensely purple permanganate (MnO 4 – ) ion

reacts with hydrogen peroxide to form the light-brown Mn2+ ion

according to the following equation:

5 H 2 o 2 (aq) + 2 mno 4 – (aq) + 6 H+(aq)

→ 2 mn2+(aq) + 8 H 2 o(l) + 5 o 2 (g)

Because the purple color of the permanganate ion is so intense,

the titrant can serve as its own indicator for this titration. As long

as hydrogen peroxide remains in excess, MnO 4 – ions are quickly

reduced to Mn2+ ions, and the solution remains a light brown

color. As soon as permanganate ions are slightly in excess, the

solution assumes a purple color. By determining the amount of

permanganate titrant required, we can calculate the amount of

hydrogen peroxide that remained in the original samples.