A Student-Designed Analytical Method

Titrations and Indicator Ranges in Mixed Aqueous-Organic Solvents

Sheryl A. Tucker and William E. Acree, ~r? University of North Texas, Denton, Texas 76203 The use of pH indicators for endpoint determination in aqueous - ti- Indicators trations provide a colorful and inexpen- 123 4 5 6 7 8 91011 12131 sive analytical method; however, many Yellow I Y-G I and bases that are used in the Yellow I Y-G I pharmaceutical industry are only spar- ingly soluble in water; e. g., lidocaine, li- Bromocresol Purple Yellow 1 Y-G I mue docaine hydrochloride, salicylic acid, Yellow I Y-G I Pur~le benzoic acid, and procaine (1). Such acids and bases can be examined in Bromophenol Blue Yellow I Y-G I Blue mixed aqueous-organic solvents, or- Yellow I Y-G I Blue ganic-surfactant systems, and in non- aqueous solvents. The question ad- Yellow I Y-G I Blue dressed in this student-designed Yellow IGI Blue laboratory experiment is, "Can common pH indicators be used in mixed solvents Congo Red Purple I Orange and if so are the color changes and pH Purple I Red ranges the same as in water?" Over the last decade, numerous acid-base titra- this Cresol Purple Yellow I Gray I Purple tions experiments have appeared in Yellow I Gray I Pur~le Journal (2-4, but to our knowledge none address this issue. Harris and Kratochvil's and Sawyer, Heineman and Cresol Red Yellow I Purple Beebe's laboratory manuals each con- Yellow I Gray I Purple tain an acid-base in a purely Yellow nonaqueous media, but both are Yellow "canned" experiments that do not allow for student design. Harris and Kratochvil's experiment, "Nonaqueous Pink 1 Red 1 Yellow Photometric Titration: Determination Pink lorangel Yellow of Oxine with p--sulfonic Acid", Methyl Yellow requires that a spectrophotometer be Red I 0 I Yellow available plus it uses 100% acetonitrile Red I 0 I Yellow which is more expensive than a mixed aqueous-organic solvent system that is Colorless I Pink no greater than 50% organic by weight Colorless I Pink (5).Sawyer, Heineman, and Beebe's ex- periment, "Acid-Base in Non- Yellow 10 1 Pink aqueous Media: Titrations of Nicotine", Yellow I Pink analyzes a toxin, nicotine, as well as using toxic glacial . Several Blue Yellow I GI Blue special safety precautions are necessary Yellow IYG/ Blue to perform this experiment; e. g., wear plastic gloves, work in a hood, explosion Thymolphthalein Colorless I Blue possibility, etc. Also the electrode can he Colorless Teal damaged if left in the glacial acetic acid for more than 20 minutes (61, and pure Tumeric Yellow-Green 101 Brown nonaqueous solvents of low dielectric Yellow I Peach constant present problems like Figure 1. Typical student data for the color changes and pH ranges of common indicators. The homoconjugation that complicate the top color bar for each indicator is for water. The bottom color bar for each indicator is for the analysis. Our experiment allows for stu- mixed aqueous-methanol solvent system (50 % methanol by weight) (Y-G or YG = yellow- dent design, is inexpensive, poses no ab- green, R = red, GR = gray. 0 = orange and G = green). normal health risks and does not require special safety precautions. It allows students to work in with the proper amount of instruction and direction by the teams to design a new analysis method to address the con- laboratory supervisor. As educators we realize the import- cerns associated with titrating sparingly soluble acids ance of teaching students not only how to use the appara- andlor bases with common pH indicators. High school teachers to college professors can employ this experiment 'Author to whom correspondence should be addressed.

Volume 71 Number 1 January 1994 71 PH nately, it is not necessary to know the Indicators 123456 7 8 9101112131 autoprotolysis constant or the pH ranges Bromocresol Green Yellow I Y-G I Blue for mixed aqueous-organic solvents as Yellow I Y-G I Blue the pH can be calculated directly via pH, = pH, - (Ex- EJS (1) Bromocresol Purple Yellow I Y-G I Blue Yellow where pH, is the pH of the sample (XI, IY-GI Purple pH, is the pH of the standard (s), Ex the millivolt value of the sample as read by Bromophenol Blue Yellow I Y-G I Blue the pH meter, E, the millivolt value of Yellow IY-GI Purple the standard as read bv the DH meter. and S the slope, normafly 59.i6 ~VI~H Bmmothymol Blue Yellow I Y-G I Blue unit at 25 'C. This form of eq 1refers to Yellow I Y-G I Blue a pH meter equipped with a combination glass electrode (8-10). The millivolt set- Congo Red Purple I Orange ting is used because a greater than the Purple I R I Orange normal pH range may be encountered. Many digital meters are capable of dis- Cresol Purple Yellow I Gray I Purple playing -19.99 to 19.99 pH. In other me- L Ydnw Purple ters, such as analog meters, the scale is limited. The millivolt range of most pH Cresol Red Purple meters provides a greater range, such as F I"r7 Purple 1400 mV, then the pH range of 0-14 that is 414 mV at 25 'C. The pH value can be Methyl Orange Red I Orange I Yellow calculated comparing the millivolt val- Orange I Light Orange ues obtained in the standards with the millivolt values obtained in a sample (8). Methyl Red Pink lRed I Yellow Standard pH values in mixed aqueous- Pink I Orange I Yellow organic solvents of varying composition are readily available in the literature Methyl Yellow Red I 0 I Yellow (11, 12). For example, a 0.05 m potas- Red I 0 I Yellow sium hydrogen phthalate reference in a mixed aqueous- Phenolphthalein Colorle~is I Pink methanol system (50% methanol by ..I"^" I mkb COIL,,,,, I r ,, ,n weight) has a pH of 5.131 at 25 'C (11).

Phenol Red Yellow 10 1 Pink Experimental Measurements Yellow I 01 Pink Students or instructors can prepare the pH indicators according to the CRC Thymol Blue Yellow I Green I Blue Handbook of Chemistry and Physics Yellow IG I Blue (13).We suggest that the students work in pairs to facilitate discussion when Thymolphthalein Colorless I Blue problems arise and that two lab periods Colorless Teal be allotted for this experiment. The aqueous ranee of the DH indicators can Turneric Yellow-Green 101 Brown be found in the "CRC;, or the students Yellow I Orange can use a pH meter equipped with a com- bination glass electrode that has been Figure 2. Typical student data for the color changes and pH ranges of common indicators. The calibrated with a standard buffer solu- top color bar for each indicator is for water. The bonom color bar for each indicator is for the tion to find the ranges via titration with mixed aqueous-acetonitr~lesolvent system (30% acetonitrile by weight) (Y-G or YG =yellow- 0,01 M (HC1) andior green, 0 =orange and G =green). 0.01 M or potassium hydroxide (NaOH or KOHL We report the use of a tus and how to follow procedures, but also how to approach mixed aqueous-methanol system (50% methanol by "real-world" problems utilizing the knowledge they have weight) and a mixed aqueous-acetonitrile system (30% learned in the classroom. This lab serves that Puwse be- acetonitrile by weight). For the mixed aqueous-methanol cause it uses the ideas that are familiar even at the high system a 0.05 molal potassium hydrogen phthalate (KHP) school level. It serves as a learning twl because it is de- buffer solution should be prepared in the mixed solvent. It signed to ensure success and build confidence in one's abil- will have a pH = 5.131 at 25 'C (11).For the mixed aque- ity to design an approach to solve "real-world" problems. ous-acetonitrile system a 0.05 molal KHP buffer solution Acid-base titrations provide a convenient means to ana- should be prepared in the mixed solvent. It will have a pH lyze many unknown mixtures. Water has an autoprotolysis = 4.985 at 25 'C (12). There are several other standards for constant ofK, = 10-l4 and hence a pH range from &14, but various mixed aqueous-organic solvent systems listed in nonaqueous solvents or mixed aqueous-organic solvents the literature (12). After taking the KHP standard's read- have pH ranges that may be larger or smaller than the ing in millivolts, the pH ranges and color changes in the normal 0-14 range. For example, the pH range of metha- mixed aqueous-organic solvents can be found by utilizing a no1 with an autoprotolysis constant of KAO-'~." is 0-16.7 pH meter. The millivolt readings can be converted easily to and acetonitrile has a K,=10-28.6and a pH range of 0-28.6 pH via eq 1. When finding the mixed aqueous-organic sol- (7). Various percent mixtures of waterlmethanol would vent pH range and color changes one must remember to have pH ranges somewhere in between 0-16.7. Fortu- prepare the 0.01 M HCI a@ the 0.01 M NaOH or KOH

72 Journal of Chemical Education A AA A BLUE 7 Ah PURPLE A tA

A GRAY

Figure 3. Titration of salicylic acid with potassium hydroxide in the Figure 5. Titration of lidocaine with hydrochloric acid in the mixed mixed aqueous-methanol solvent system (50% methanol by weight). aqueous-methanol solvent system (50% methanol by weight). The The different symbols represent the color changes for the pH indica- differentsymbols represent the color changes for the pH indicator tor mcresol purple (squares = yellow, circles =gray and triangles = bromocresol green ( squares = yellow, circles = yellow-green and purple). triangles = blue).

*.. *.. ORANQE .

...... YELLOW

5 10 15 a A 30 d5 mL1 KOH

Figure 4. Titration of otoluic acid with potassium hydroxide in the Figure 6. Titration of lidocaine hydrochloride with potassium hydrox- mixed aqueous-acetonitrile solvent system (30% acetonitrile by ide in various solvent systems. The circles represent a two-phase weight). The different symbols represent thecolor changesforthe pH surfactant system of aqueous-organic cetylpyridinium chloride. The indicator tumeric (squares = yellow and circles = orange). squares represent a mixed aqueous-methanol solvent system (50% methanol by weight). The triangles represent an equal volume of solutions in the mixed aqueous-organic solvent of interest. dichloromethane and the surfactantcetyipyridlnium chloride. Each gmup of students should prepare a 0.01 M solution of a sparingly water soluble acid and base for titration in the Discussion mixed solvent system being examined. Acids such as sali- Typical student data illustrating the pH indicators' color cylic acid, o-toluic acid, benzoic acid or lidocaine hydrochlo- changes and pH ranges in water and in themixed aqueous- ride will work, and bases as lidocaine, sodium benzoate or methanol system are shown in Figure 1. Note that the sodium salicylate (those acids listed here will go to their indicators' pH ranges andlor color changes are rarely the basic form by the addition of excess base and vice versa). same in the mixed aqueous-organic solvents. A similar Students should plot the titration curve and choose one or chart could be prepared by the student to aid in their indi- two appropriate indicators and cany out the titration in cator selection for the acid and base samples. Also note them. A plot of the titration curve will demonstrate that some transition colors were not apparent in both sol- vents. For example, the tumeric indicator has an aqueous whether they have chosen wisely. According to Williams color transition of yellow-meen to orange to brown, hut in and Howell (14) visual enhancement of titration endpoints the mixed aqueous-methanol system only a yellow to pale can be achieved by carrying out the titrations in polysty- oranee- transition is seen. The indicator m-cresol red rene foam coffee cups. We have found that the mixed aque- illustrates that we all perceive colors differently as most ous-organic solvents do not affect the integrity of the cups. students thought the basic region of the m-cresol red indi- We recommend them because they are inexpensive, they cator was more purple than red. Figure 2 is typical data for ~rotectthe electrode from accidental damaee.-, and thev can the pH indicators in water and in the mixed aqueous-ace- be reused upon rinsing. More advanced students also can tonitrile system. Once again the color changes and pH be asked to calculate the titration error involved with their ranees- differ in the two solvents. Com~arine- Fieure - 1and choices of indicators (15). Figure 2 we also note that the two mixed aqueous-organic

Volume 71 Number 1 January 1994 73 systems differ on their pH ranges and color changes as ex- not only learn how to approach problems, but also they will pected. Figure 3 is the titration curve of salicylic acid ti- gain exposure tn a littl,.:knownares ot'chemi~try~hdrma- trated with KOH in the mixed aqueous-methanol system. crutics. They will learn h~~\vto test their hypothcsis that

The color changes and pH ranges for m-cresol purple are DH indicatnrs do do not chmee their ranecs und rolor tran--- depicted. This indicator choice was appropriate because sitions in solvent systems o&er than water, and they will the color change occurs at or near the "true" equivalence have used that knowledge to approach the problem of point. The titration curve of o-toluic acid titrated with acid-base titrations of sparingly water soluble samples. KOH in the mixed aqueous-acetonitrile system with tume- Also a note for the instructor, the students will not be able ric as the indicator, and that of lidocainetltrated with hy- to merely look up the pH indicator ranges and color drochloric acid in the aqueous-methanol system wlth changes for the mixed aqueous-organic solvent systems as bmmocresol green as the indicator are shown in Figures 4 in the case of aqueous systems because they are not cur- and 5, respectively. Lidocaine or lidocaine hydrochloride rently available in the literature. cannot be titrated in water because it is not soluble and will crystallize out of solution. Figure 6 shows the im- Acknowledgment movements that can be made in the titration curves bv The authors would like to thank Glenn H. Brown for pro- ;hanging the solvent system. The titration of lidocaine hi- viding the pH indicators. drochloride in the mixed aqueous-methanol system has slightly improved the titration cuwe and hence endpoint Literature Cited detection over the two-~hasesurfactant svstem of aoue- 1. Jahansson, P -A; Hoffmaffma, G.; Stefanaaon, U.Ano1. Chim. Ado 1982,140,7748, ous-organic cetylpyridi~iumchloride, and the water sblu- 2. Harvey. D. T J. Cham. Educ 1981.68.329331. 3. Flair. M. N.;WiUiams, N. S. J Chem. Edue. 1990.67,795-796. bility pmblems have been eliminated in both cases. The 4. McClendon. M. J. Chem. Edvc 1984.61. 1022-1024. titration lidocaine hydrochloride in equal volumes of 5. Hams, W E.; Kratoehvil. B. Chemiml Sepomtiona and Mmsummmts; Saunders College Publishing: Philadelphia, 1974, pp 15b164. dichloromethane and the surfactant cetylpyridinium 6. Saver, D. T.: Heineman, W. R.; Beebe, J. M. Chemistry E~p~rimenlsfiiIn~frumfn- chloride is clearly the best in this case (see Ref. 16 for ex- la1 Methods: John Wiley 8 Sons: New York, 1984, pp 3642. perimental details). Another interesting approach to this 7. Fritz. J. S.Acl&Bose Titrations in No~g.ueourSolvents: Myn and Bacon: Boaton, 1973. o 35. experiment would be to examine the edible acidhase indicators' color changes and pH ranges, such as red cab- bage and aaDe ",iuice. in the mixed aoueous-oreanic solvent syitems.~ibaneand Rybolt rep& their :aqueous pH ranges and color chanees as well as the ~re~arativemeth- 11. Mussini T.; Covington, A. K, Dal Pozzo, F: Longhi, P; Rondinini, S.Ann. Chim. . . 1982, 72(3-4),20%211. odsfor several common and not so common edible pH indi- 12. L0nghi.P; Musaini,T.;Rondinini, S.ilnol. Cham. 1986.58,2290-2292. cators (I7). 13. Linde, D. R., Ed. CRC Handbook of Chemistry and Physrcs; CRC Press: Boston, 1990.71. pp%12. This laboratow ex~erimentis inex~ensiveand uses vir- 14. WiUiams, H.P.; Howell, J. E. J. Cham. Educ. 1988.66.680. tually no equipmeniother than a biret and a pH meter, 15. Skoog, D. A; West, D. M.; Holler, F. J Fundament& o/Anolylicol ChamiaVy: 6th ed.; Saunders College Publishing: New York. 1992, p 235. and it can be ~erformedat almost any level with the cor- 16. Rcker, S. A; Amszi, V. L.;Aeree, Jr. W E. J. Chem. Educ 1993.70 &82. rect amount of direction from the instzkctor. Students will 17. Mohane, R. C.; Ryblt, T R. J Cham. Educ. 1986,62,285.

74 Journal of Chemical Education