In the Classroom
edited by
JCE DigiDemos: Tested Demonstrations Ed Vitz
Kutztown University
Kutztown, PA 19530
Phenolphthalein–Pink Tornado Demonstration
submitted by:
checked by:
Bruce R. Prall
Math and Natural Sciences Division, Marian College, Fond du Lac, WI 54935; [email protected]
Jeannine Eddleton
Department of Chemistry, Virginia Polytechnic Institute and State University, Blacksburg, VA 24061
The titration of HCl with NaOH has traditionally been
used to introduce beginning chemistry students to the concepts
of acid–base chemistry and stoichiometry. This demonstration
utilizes this reaction as a means of providing students an opportunity to observe the dynamic motion associated with a swirling
vortex and its effect on the reaction.
Materials
1000 mL of 0.01 M HCl
10 mL of 50% (w/w) aqueous NaOH
(stored in a polyethylene container)
Phenolphthalein indicator solution
Magnetic stirrer with stirring bar (3/8 × 1½ in.)
2 Beral droppers
White backdrop
1 L graduated cylinder
Procedure
Fill a 1 L graduated cylinder with 0.01 M HCl to within
2–3 cm of the top of the cylinder. After placing a magnetic stirring bar in the graduate cylinder, place the cylinder on a magnetic stirrer. Adjust the stirring rate to form a smooth non-turbulent
vortex. Place the white background behind the graduated cylinder. Add 3–4 drops of phenolphthalein to the center of the HCl
solution near the top of the graduated cylinder. Immediately
after adding the phenolphthalein, slowly add drop-by-drop 3–5
drops of 50% (w/w) NaOH solution in the center of the graduated cylinder. If the NaOH is added too quickly the entire top
of the cylinder tends to turn pink and the vortex pattern is not
clear in the top part of the cylinder. After the initial additions,
one observes a pinkish–red shifting tornado-like column that
subsequently disappears. There may be a time lapse before the
vortex becomes visible. Allow time for observations before repeating the procedure, which can be accomplished by adding the
phenolphthalein and NaOH as previously described. If sufficient
phenolphthalein is added the first time, only the NaOH needs
to be added for subsequent observations. Additions of NaOH
can only be repeated a few times before the entire solution turns
pink. However, the solution can be “recharged” by addition of
concentrated HCl if a greater number of additions of NaOH
(more pink vortexes) are desired.
Hazards
An apron, safety goggles, and latex gloves should be worn
during the preparation and performance of the demonstration.
Concentrated hydrochloric acid used to prepare the 0.01 M
solution is extremely caustic and should be handled with care.
Sodium hydroxide and the concentrated solution used in this
demonstration are also extremely caustic and corrosive. Any
skin or eye contact with the aforementioned chemicals requires
immediate flushing with water. Any accidental spillage should
be cleaned up immediately.
Discussion
The phenolphthalein–pink tornado demonstration not
only provides students the opportunity to observe the swirling
tornado-like motion of a vortex (1), but also provides an instructor with an excellent opportunity to introduce the concepts of
limiting reagent (2), reagent in excess, as well as Le Châtelier’s
principle. The reactions relevant to this discussion are
OH (aq) H3O (aq)
H Ind(aq) 2H O(l)
2
2
colorless
2H O(aq) Ind2(aq)
3
(2)
(pH 8 10)
pinkish–red
Ind2(aq) OH (aq)
pinkish–red
(1)
2H2O(l)
IndOH3(aq)
(3)
colorless
(pH 11)
where H2Ind is phenolphthalein.
Initially, upon adding the phenolphthalein and the NaOH
to the vortex of the HCl solution the OH− concentration is
greater than the H3O+ concentration in this localized region
of the solution, because the OH− does not sample all the H3O+
in the solution. This incomplete mixing results in the vortex
becoming pinkish–red, which is what one would predict to
happen based on applying Le Châtelier’s principle to reaction 2.
The decrease in H3O+ concentration causes a shift to the right
in reaction 2 and correspondingly an increase in the Ind2− concentration. The continued swirling action of the vortex tends
to draw the OH− to the bottom of the graduated cylinder and
the OH− eventually reacts with the H3O+ at the reaction front.
Because H3O+ is the reagent in excess, after all the added OH−
has been consumed, the H3O+ concentration becomes greater
than the OH− concentration at the reaction front and causes
the equilibrium for reaction 2 to shift to the left rendering the
solution colorless. However, as additional NaOH solution
is added, the OH− concentration becomes greater than the
H3O+ concentration, which is evident by the pinkish–red color
remaining in the solution.
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 4 April 2008 • Journal of Chemical Education
527
In the Classroom
Since the H3O+ + OH− reaction is extremely fast, reaction
3 offers the best explanation for the persistence of the pinkish
color within the vortex as the pH drops. In Nicholson’s kinetic
study of reaction 3 (3), she reported that the reaction is relatively
slow in 0.05 M NaOH solutions (pH 12.7), but is quite rapid
in 0.30 M NaOH solutions (pH 13.5). If one assumes an initial
vortex column of 10.0 cm in length with a radius of 0.50 cm, the
pH of this volume (7.85 mL) following the addition of 4 drops
(0.20 mL) of 50% (w/w) NaOH (approximately 19 M) can be
calculated to be 13.7. This calculation takes into consideration
the loss of NaOH owing to its reaction with 7.85 mL of 0.01
M HCl. Since reaction 3 is very fast at this pH, it provides a
pathway for storing Ind2− as IndOH3−. This could account for
the pinkish–red color in the vortex subsisting over a few seconds.
The radius of the visible vortex slowly decreases, resulting in the
visible vortex eventually disappearing as the pH falls below 8.
An additional reason for the persistence of the pinkish–red
color could be the localization of the OH− ion within the vortex.
This limits the reaction to taking place only at the interface with
the bulk solution, making it slow compared to the rate in a homogenous mixture. Modifications of the demonstration include
replacing the 0.01 M HCl solution with either 0.1 M HCl or
0.01 M HC2H3O2. For larger classrooms it may be appropriate
to scale up the procedure by a factor of four, that is, replace the
1 L graduated cylinder with a 4 L beaker.
Literature Cited
1. Shakhashiri, B. Z. Chemical Demonstrations; University of Wisconsin Press: Madison, WI, 1983; Vol. 1, pp 271–279.
2. Hill, J. W.; Petrucci, R. H. General Chemistry, An Integrated
Approach; Prentice-Hall: Upper Saddle River, NJ, 2002; pp
109–112.
3. Nicholson, L. J. Chem. Educ. 1989, 66, 725–726.
Editor’s Note: Ed Vitz
The pink tornado may seem paradoxical in its persistence,
given that the reactions that cause dissipation of the color are
among the fastest known (1, 2). This demonstration contrasts
the rapid reaction with stirring in a circular direction to form the
pink tornado, with the slow radial migration of ions that causes
dissipation of the pink tornado. Why doesn’t the color dissipate
quickly, reflecting the large rate constant for the reaction of
OH− with H3O+? Proton transfers are “diffusion controlled”
because they occur at virtually each encounter, have zero activation energy, and thus occur at rates that depend only on the
rate at which the reactants approach each other. Not only are
proton transfer reactions diffusion controlled, but their rates
are enhanced by processes that lead to apparent translation
without reactant migration. For example, the proton exchange
H3O+ + H2O → H2O + H3O+ is equivalent to the water
molecule moving into the place of the hydronium ion and vice
versa (this is implicit in the Grotthus mechanism). The apparent translation is proportionately greater for larger aggregates
of water molecules, with four or more molecules likely to be
involved, and in ice, where the hydrogen-bonded structure is
more permanent, the observed rate constant is 1013–1014 M‒1
s‒1 (3). In these cases, when a proton moves by a distance approximately equivalent to bond length, the H3O+ is apparently
528
translated by a distance equal to the diameters of several molecules. Why, in spite of this very rapid reaction, does the color of
phenolphthalein remain in the vortex and not move out radially
from it? In the radial direction there is no mixing by stirring, and
the reaction front (the boundary between the red, basic solution
and the colorless acidic solution) barely seems to move, just as is
observed in a totally unstirred mixture.
In spite of the large rate constants involved, diffusion leads
to very little mixing (or progression of a colored reaction front,
as we see in the pink tornado) (2). Convection and mechanical
stirring are necessary to make even the fastest reactions proceed
at convenient time scales! Even when mechanisms for translation without migration are included, diffusion is enhanced only
slightly (4). As a matter of fact, the rate constants of reactions
that lead to augmentation of diffusion by a reaction (like electron exchange and presumably proton exchange) can actually
be determined from diffusion coefficient measurements (4).
Amazingly, the rate of diffusion is only enhanced measurably
by the electron exchange mechanism if k > ~108 M‒1 s‒1 under
typical conditions. Direct measurement of the diffusion coefficient has been used to measure the rapid ferrocene–ferricene
electron exchange rate constant (2 × 1011 M‒1 s‒1) (4–6). What
one observes in the fastest diffusion-controlled reactions is very
slow progression of a reaction front, such as the persistence of
the red of phenolphthalein in the pink tornado, in spite of the
“rapid” reaction with the surrounding solution.
Thus it may not be necessary to invoke the formation and
decay of species such as InOH3− to explain the persistence of the
pink tornado, although this species is probably formed to some
extent. It is interesting to add a drop of phenolphthalein to ~20
mL of stirred 1 M NaOH, wait for the color to fade as InOH3−
is formed, then add HCl as stirring continues. The pink color
appears rapidly and disappears just as rapidly as the equilibria
are established with stirring. So it may not be the “slowness” of
the reaction of InOH3− with HCl, but merely the “slowness” of
radial diffusion of H3O+ to the tornado, which accounts for the
persistence of the pink.
The pink tornado demonstration is excellent for contrasting
the relatively rapid rate of reactions in mixed solutions with the
paradoxically “slow” rates of the same reaction where mixing is
ineffective.
Literature Cited
1. Caldin, E. F. The Mechanisms of Fast Reactions in Solution; IOS
Press: Amsterdam, 2001; pp 227–263.
2. Jordan, P. Chemical Kinetics and Transport; Springer: New York,
1979.
3. Caldin, E. F. The Mechanisms of Fast Reactions in Solution; IOS
Press: Amsterdam, 2001; p 233.
4. Ruff, I.; Friedrich, V. J. J. Phys. Chem. 1971, 75, 3297–3302.
5. Ruff, I.; Korosi-Odor, I. Inorg. Chem. 1970, 9, 186–188.
6. Ruff, I.; Friedrich, V. J.; Demeter, K.; Csillag, K. J. Phys. Chem.
1971, 75, 3303–3309.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2008/Apr/abs527.html
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Journal of Chemical Education • Vol. 85 No. 4 April 2008 • www.JCE.DivCHED.org • © Division of Chemical Education