edited bv ROBERT REEVES Marlborough Schwl 250 S. Rossmore Avenue LOS Angeles. CA 90004
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Kinetics of the Fading of Phenolphthalein
Lo$ Nicholson West Springfield High School, Fairfax County. VA 22152
Phenolphthalein is used in several diverse applications. Chemistry students may he aware that phenolphthalein is an active ingredient in some laxatives, hut they most frequently encounter phenolphthalein as an acid-base indicator to determine the equivalence point in a titration. If excess hase is oresent in the flask a t the end of the titration, thestudent may note that the pink color of phenolphthalein fades if the mixturestands for a while. But thisslow iadina is of no consequence to the titration, and the solution is discarded without further thought. Yet this fading of phenolphthalein in alkaline solution is interesting in its own right and can serve as the basis for an experiment illustrating pseudo-first-order kinetics. The procedure is extremely simple and foolproof, uses common reagents and equipment, and gives excellent results. Slruciural Forms ot Phenolphthaieln Although phenolphthalein is one of the most common acid-base indicators, its chemistry is not that of a simple conjugate acid-base pair, HIn-In-. The structures of the important forms of phenolphthalein are shown in Figure 1. Phenolphthalein is colorless a t pH 8 and lower. This colorless form has structure 1, abbreviated HzP. As the pH rises from 8 to 10, both phenolic protons are removed with approximately equal ease and the lactone ring opens, producing the familiar red-pink form with structure2,abhreviated Pi-.At still higher pH, the pink color slowly fades, producine structure 3. ahhreviated POH1-.' All color chanees are reversible, a n d while the conversion of H2P to p2--is extremelvraoid and essentiallvcom~letebv the time the DHis as highas il,the conversioiof ~ 2 to ' ~ 0 ~ at3 higher is sufficientlv slow that its rate mav he easilv measured. Since P2- is intensely colored, this con;ersion ofP2- t o POH3- can he monitored hy measuring changes in absorbance, or optical density, of an alkaline solution of phenolphthalein.
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Determination of the Pseudo-Order in Phenolphthalein
The fading of phenolphthalein in hasic solution can be represented by the reaction: P2- OHPOH3-, and the rate law can he expressed rate = k[OH-]"'[P2-I". However, our orocedure uses stronelv basic solutions containina 0nlv a trace of phenolphthalei< so the OH- concentration exceids that of ohenolohthalein hv a factor of at least 10' in any mixture: ~here'fore,duringkach run, the OH- concentration remains essentially constant, and the rate law then becomes rate = kl[P2-1". In this new rate law, kl = k[OH-I"', and the
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' Lalanne, J. R. J. Chem. Educ. 1971,48,266-268.
2Masood, A.; Shastri, N. K.; Krishna, B. Chim. Anal. 1970, 52, 1289-1295.
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3 Figure 1. The sbuctural forms of phenolphthalein.
reaction is said to he "pseudo nth order" in phenolphthalein (or in soecies P2-). If the reaction is pseudo first order in phenolihthalein (or n = I), a plot of-ln [P2-] versus time should give a straight line with slope equal t o -kl. According to Beer's law, the spectral absorbance of the solution is directly proportional to [P2-1, or [P2-] = constant X (ahsorhance), so ln[P2-] = In (constant) In (ahsorhance). Therefore, a plot of in (absorbance) versus time should also be a straight line with slope of - k l , if the reaction is pseudo first order in phenolphthalein.
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Experimental Solutions and Equipment NaOH solutions in the range 0.05-0.30 M give convenient rates of fadingof phenolphthalein.For a particular concentration of NaOH, the rate of fading of phenolphthaleinincreases with increasing ionic ~trength.~ This dependence of rate on ionic strength can be explained by the fact that the reaction involves the approach of two negatively charged ions, and their mutual repulsion is decreased in Volume 66 Number 9
September 1989
725
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Flgvre 2. Ln (abmbance)versus time. each at wnstant [OH-]. In all graphs. the points are data, and each line is drawn using least-squares analysis.
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