edited by ROBERT REEVES Marlborough School 107 South Wilson Drive
Preparation of Potassium Hydrogen Thomas Rees Phillips Academy, Andover. MA 01610'
Potassium hydrogen maleate (KHM) is an acid salt useful to have in the stockroom. I t can be used as an "unknown" acid in demonstrating the strength of weak acids. Unlike most common acids, a solution of KHM turns methyl orange a deep orange without a trace of red. A moderately soluble salt with a molar mass of 154.16 g, KHM can also be used as an "unknown" acid in acidimetric titrations. KHM is easily prepared by students in two simple steps, beginning with maleic acid
Because potassium hydrogen maleate is fairly soluble in water the solutions must be rather concentrated. and the final solution must be chilled in an ice bath. If sodium hydroxide is used instead of potassium hydroxide, the resulting acid salt is less soluble, but unfortunately sodium hydrogen maleate exists as the trihydrate2, introducing an undesirable complication to the experiment. This preparation of potassium hvdroaen maleate can be used a t almost any time in the school year after the students have been introduced to the mole concept, equation balancing, molarity of solutions, and stoichiometric calculations. If the experiment is performed early in the year, the experiment gives a practical example of the necessity of stoichiometric calculations. The studen{s also have the pleasure of seeine a crvstallization. a new exnerience for most students in thiir first chemistry'course. B; the use of indicators the students discover that maleic acid is a stronger acid than the hydrogen maleate ion. Finally, the students learn the importance of observing and reasoning because the second equation is not given to the students, and they must deduce the formula of the crystalline product. If the experiment is performed later in the year and the students already know the Bronsted theory of acids and bases, the experiment demonstrates that the X2-ion is sufficiently basic to remove aproton from the H2X molecule. The instructor may also choose t o demonstrate the titration of KHM with perchloric acid in glacial acetic acid, showing maleate (HM-) ion can act as a base when that the hvdroeen . reacting with a strong enough acid. Actually, students can perform this titration. but the use of laree amounts of elacial acrtir acid is expe~~sive and van be disaireeahle. If the students have studied some oraanic rhernistrv. tbev can be shown that the structure of thekaleic acid miiecuie
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Present address: 35 Sunset Road, Hamden, CT 06514. 2Temple,J. M. J. Amer. Chem. Soc. 1929, 51, 1754. Weiss, J. M.; Downs, C. R. J. Amer. Chem. Soc. 1923, 45,2341. Hunter, L. Chem. Indos. 1953, 155. Shahat. M. Acta Cryst. 1952, 5, 763. Westheimer, F. H.; Benfey, 0.T. J. Amer. Chem. Soc. 1956, 78. 5309.
gives an explanation for the great contrast in acid strengths of maleic acid and the hydrogen maleate ion.3
The internal hydrogen bonding attracts electrons from the upper carboxyl group, as shown above, thus promoting the release of H+ from that carboxyl group. The dissociation is large for a carboxyl group. The constant, K , = 1.14 X and this HM- ion is a much weaker acid, Kz = 5.95 X weakness can be attributed to the internal hydrogen bond, which tends to hold the H+ within the ion. If an alert student objects to the seven-membered ring, the teacher can point .O distance is even less than in the out that the O-H. formic acid dimer. When the vials of KHM are collected by the instructor, the samples that gave a good titrimetric analysis are worth saving. For reasons already mentioned this compound is a useful addition to the stockroom supply of acids.
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Notes on the Experiment Asolutiun uf6.96g msleic acid in l ~ m l w a t ris r neutralized to the p h e n o l p l ~ t h a l e i n c n ~ l pwith ~~~n 20~ mlfi.00Af KOH.'I'o this neutrnl)zed w l ~ t i mi i sclded n second portion uf 6.96 p: molcic acid in 15 ml water. Chilling in an ice bath gives white crystals of potassium hydrogen maleate. Because of the solubility of potassium hydrogen maleate, the yield is typically around 8 g or 43% of the theoretical yield of 18.5 g. The time required for the preparation of potassium hydrogen maleate is one 2-hour laboratory period, provided that the student has bad a small amount of laboratory experience. At Phillips Academy, sometimes the students have had no laboratory experience at all. These students needed two 2-hour laboratory periods, the first period being used for the neutralization of the maleic acid. However, if a student performed a poor neutralization, that procedure could be repeated and the entire preparation finished in the next 2-hour period. The experiment can he interrupted at virtually any time, hut a convenient time is after the solutions of dipotassium maleate and rnaleic acid have been mixed. After the crystals are dry, another laboratory period is needed to analyze the product by acidimetric titration. If the titrant is standardized 0.1 M NaOH, a sample in the range of 0.460.61 g will consume approximately 30 to 40 ml of the NaOH. The sample is dissolved in 15 ml distilled water and titrated to the phenolphthalein endpoint. If there is not time for the students to perform the analysis, theinstructor may prefer to give thestudents thedata for a titration, letting the students calculate the equivalent weight of the compound.
Volume 63
Number 2
February 1966
157
Post-Laboratory Actlvltles
One final activity connected with this lab is its use to develop problem-solving ability. Frequently my students are the type who demand a strong intellectual challenge and tolerate the frustrations of wrestling with a problem whose answer is not immediately clear. I ask them to propose a reasonable formula for the product of the reaction between the dipotassium maleate and the maleic acid. Although they have enough information, the students need strong guidance plus the information from the titration before they are able to make a reasonable proposal. The incomplete equation is written on the blackboard, and the students cooperate in listing the laboratory observations. H&H204
+
0.060 mole reacted with soluble in water acidic to litmus
KzCJb04
+
0.060 mole soluble in water pink to phenolpthalein
stronger acid melts, burns, leaving no ash
crystalline product less soluble acidic to litmus weaker acid chars without melting, leaving an ash. Further heating gives a pale pink-violet color to the Bunsen flame.
The students copy the data from the blackboard. In the discussion, they are told that under the gentle conditions of the preparation the C2H402 portion of the molecule is not broken apart. Not surprisingly, the Phillips Academy students have come to call this exercise "Suffering and Aeonv". I tell the students that often a scientist has toagonizeovkr a problem only to find that the answer, once gained, is surprisingly clear. The students are assured that they will not be peualized if they are slow to arrive a t the correct formula. The important thing is to write a clear report showing how they arrived at the formula. Whenever a student is ready to propose a formula, he or she reports it privately to theinst& tor, giving the reasoning behind the proposal. All of the students are told that if thev arrive a t the correct formula. they should not reveal it gut should let their classmates continue to "suffer". The students honor this request and
158
Journal of Chemical Education
enter into the spirit of the exercise with good humor because they know that their grade is not affected if thev arrive a t the . answer slowly. Although there is sufficient information t o make a reasonable proposal without titration data, no student has been able to do so. The titration gives an equivalent weight in the range 153-155; results outside that range require-a repetition. With this additional information, only one student out of 88 has failed to arrive a t the formula KHC4H20a.This student was finally told the answer and the reasoning behind it and then was required to write the normal report. I t is important to admit to the students that letting their classmates suffer is not good scientific practice, hut that communication among scientists is most important. The reason for the secrecy is to let the other students have the experience of solvingthe problem. Obviously "suffering and agony" cannot be used year after year. word-gets around. ~ h e k x p e r i m e nwas t usedmostly in the summer sessions, in which the student body was new every year. In the academic year the experiment was not used again until all the previous students had graduated. I t is possible, of course, to use this experiment as a study in acids, bases, and chemical equilibrium, omitting the feature of "suffering and agony". The Renor1
The amount of material in the report depends upon how much chemistry the student has learned before doing the experiment. The report also depends upon the way the instructor uses the experiment. The student should calculate from the experimental data the ratio of the number of moles of potassium hydroxide to the number of moles of maleic acid. The student should reoort the vield of nrodnct and the percent of the theoretical iield. If titratidn is performed, the student should show the data and calculations leading to the equivalent weight. If the student is required to deduce the formula of the product, the student should write an essay showing how the experimental observations lead to the formula. If the student has already studied the Bronsted theory, the report can include a discussion of the acid strengths of HZmaleate and HM- and the base strengths . of HM- and maleate2-. The laboratory handout used with this experiment is available from the Office of High School Chemistry, American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036. Ask for experiment 1003.
a
