A. G. Spliwgerber, D. 6. MacLean, and J. Neils Gustavus Adolphus College st. Peter, Minnesota 56082
A Unified Introductory Chemistry Laboratory
There are several objections to conventional introductory chemistry laboratories, some of which apply to the laboratory portion of advanced courses as well. These objections center more on how chemical principles and operations are to be presented rather than on which of these principles and operations properly belong in the introductory laboratory. A student usually works on a new experiment each laboratory period; these experiments are generally unrelated. This leads (1) to an unnatural chemical situation which does not approximate closely enough a real research experience. There is generally a fixed amount of time for a given experiment. Some students consequently feel rushed while others consistently finish quickly. There is no easy way to adapt the varied working paces of the students to the experiment. The results of a student's work are generally too easily predictable and the materials he works with too artificial ( f ) . He spends his time redetermining values which are too easily found in chemistry handbooks and other easily available reference sources. It is not much fun for the student to work on something he knows thousands of other students have done in the past. There is therefore a natural temptation to "dry lab" results. One answer is to let students work on projects independently. The objection to this, apart from large enrollments, is that most students are not equipped to work completely independently (3). With students of widely differing chemistry backgrounds, it is inevitable that some of them will become thoroughly lost and discouraged by this approach. Another approach (1) is to devise a project-type unified laboratory program which better parallels actual research procedures than a series of disconnected experiments. The difficulty in such a program is to expose the beginning student to as many as possible of the techniques, manipulations, and laboratory operations which should be found in an introductory laboratory course (b,4). At Gustavus Adolphus College we have developed a laboratory program which meets some of the objections cited. This program was initiated with a small group of students, and will be used for the fall term which runs approximately 14 laboratory periods. Enough material is present that a longer period of time may easily be taken.
equivalent weight are determined. The ionization constant of the acid is determined, followed by preparation of the copper salt. This is analyzed for waters of hydration and percent copper. Copper concentration of a saturated salt solution is determined by spectrophotometric and gravimetric means, and values for the solubility product constant calculated. Silver salts may be prepared by similar means and solubility products determined by gravimetric, titrimetric, and potentiometric measurements. Synthesis of Benzoic Acids One of seven related benzoic acids is synthesizedfrom a possible eight starting compounds issued as unknowns (Table 1). This is conveniently accomplished (6) by refluxing 3 g of starting material with 9 g KMnO, in 300 ml of water in a 500-ml flask f0r.a. period of 2-3 hr. Filtration of the MnOl precipitate followed by acidification of the filtrate yields the crude acid which is filtered and allowed to dry on a wstch glass. Exceptions occur in the cases of o-anisie and o-nitrobensoic acid, which are obtained by evaporation of the filtrate after acidification. Yields (Table 1) are high except for the more soluble acids (benzoic and ochlorobeneoic).
Table 1.
Starting Material
nenssldehyde nenzy~ AIW~OI
o-Nitrotolueqe p-Nitrotoluene o.Chiorobenrsldehyde p-Chiorobenzddehyde 0-Aniaaldehyde p-Aniaaidehyde
Yields and Conditions
Acid
Denmic B ~ ~ S O ~ O
o-Nitrobanroic p-Nitrobenroic o-Chiorobenzoic p-Chlqrobenzoic a-Anlso pAnisic
Crude rar Aoid tion
Conditions
44 43 98
50 50 62
Water
76 56
98
Whter I : 1 Hn0:EtOH Water 2:1 Hz0:EtOH
83 70 84
58
72 58
97
water Water
2:1H10:Et0H
The student identifies his unknown acid in subsequent experiments, and when this point is reached he may celculate percent yield of crude acid. At this time the pertinent oxidationreduction equation may be balanced as an exercise.
Purification The student next purifies his crude acid by recrystallization. Experimentation is required to determine whether water alone is sufficient as a. solvent or a, mixed solvent (water-ethanol) necessary. A good apprortch is to first try 50 ml of water per gram of crude acid, and then add more water, ethanol, or bath according to judgment. Typical yields and solvent conditions used me also listed in Table 1. I n subsequent experiments the student will run out of purified acid. A aample of commercial purified acid may be made available to him a t this time.
The Laboratory Progmm
Molecular Weight and Equivalent Weight
The students begin with the synthesis of benzoic acid or one of its monosubstituted relatives. The material is purified, and its melting point, molecular weight, and
The purified material k next used to determine moleculer weight by freezing point depression and equivalent weight by titration of the acid with standard base. For the freezing paint determinations, acetamide has been found to be a. goad solvent
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Jourrml of Chemicol Educofion
Table 2.
Molecular and Equivalent Weights Molecular Weight Freezing Freezinz point. Point
Acid
Equivalent Weight Weieht Titration
Table 5.
Solubilitv Products Determinations. C o ~ ~ Salts er
Molaoular Weight Weinht Aotual
Table 6.
Water and Copper Analyses
Table 3. Acid Melting Points Melting point experimental
Acid
122-130 10~107 184-193 134-163 242-254 142-154
Benroic o-Anisic p-Anisic o-Chlorobenzoic p-Chiorobenzpic o-Nitrobenaola rrNitrobenaoic
Bsnaoic o-Aniaia p-Anisio o-Chlorobenroio p-Chlorobensoia o-Nitrobeneoio p-Nitrobenzoic
122.4
4.0 3.91
5.77 2.82 5.00 2.25
4.23
P K ~ Literature 4.2
4.09 4.47 2.92 3.98 2.17 3.42
~xperi- thee mental retioal anh. salt
salt
etry
dketes
142 243 146
241
Eauilibrium Constants
PK-. Expenmental
~
copper eslt
100-101 184-186
242-246
Table 4. Acid
Melting point literature
Table 7. Titration Conditiona
Solubility Product Determinations, Silver Salts
silver asli
K-
Gravlmetric
K'. Titrimetrio
K..
Potentlometric
Water Water 3 : l Et0H:H.O Water 4 : l Et0H:H.O Water 3:l EtOH:HzO
in which the acids do not dimerize. A Kf of 4.20 is a reasonable value for this group of acids, and this was used to determine the moleculer weights recorded in Table 2. These vdues were determined using an ordinary 1lO0C thermometer. Better acmmcy may be had using a. thermometer with 0.1 degree divisions if available. One and a half grams of purified acid in 15 g acetamide allows a convenient volume for theusual test tubeapparatus. As indicated in Table 2, accuracy to within 10% of the true value can be expected. Equivalent weight measurement may be carried out by titration of 0.5 g amounts of purified acid with standardized 0.1M base. The student may also standardize his base using potassium acid phthalste or oxslic acid. Typical vdues are listed in Table 2, and are seen to lie within 2% of the true molecular weight in all cases. Comparison of freezing point with titration results allow the student to estimate the number of replaceable hydrogen atoms in his acid. These values dso provide a preliminary identification of his acid, provided a list of possibilities is made available.
Identification of the Unknown Acid Melting points and mixture melting points are taken next in order to make a positive identification. Experimental values and typical melting point ranges are shown in Table 3 together with literature values (6). Examination of these values shows that the melting point alone cannot be used in all cases to identify the acid; a. combination of equivalent weight determination and melting point is required. A mixture melting point using availd e samples of known purified acids may be used to complete the identification.
Equilibrium Constant Equilibrium constants are determined by finding the pH at half neutrdizetian with standard base. The values shown in Table 4 were determined using a. standardized Corning Model 7 pH meter. Comparison of these apparent pK's with thermodynamic pK values from the literature (7) shows good agreement in the case of the more water soluble acids and poorer agreement where determinations were made in ethanol-water mixed solvents. If desired, more extensive experiments may be performed (8).
Salts Sodium s d t s are readily made hy neutralization of the acids with sodium hydroxide. Yields are essentially 100% in a11
cases. These may he used to prepare a. variety of less soluble salts, of which. the copper and silver salts were chosen for further study.
Copper Salts, Solubility Measurements To prepare the copper salts, approximately 1 g of the sodium salt in water is mixed with a stoichiometric amount of C U ( N O ~ ) ~ . 3H10. Total liquid volume is fixed a t 50 ml except in the case of the ortho-nitro derivative, where 25 ml is better. The precipitated copper salt is filtered and dried and the supernatant saved for andysis. The supernstsnt is treated with excess ammonia, and copper concentrations are measured using the 600 mp absorption of the copper-ammonia complex ion ( 9 ) . This exercise requires that copper standmds be made up m d a Beer's Lsw plot constructed (10). Assuming that $1 copper in solution is in the form Cua+, s value for the solubility produd of the copper salt may be calculated. Results are shown in Table 5. A Beusch & Lamb Spectronic 20 was used for all colorimetricmeasurements. Once the molecular formula. of the hydrated copper salt is known (the student may determine it), the weight of dried precipitated copper salt may be used to calculate a second value for the solubility product. These are also listed in Table 5. The grrtvimetric values given take waters of hydration into account, and are seen to compare well with the speetrophotometric determination. It should be noted that whet the student measures is not really a solubility product. With first stability constants being of the order (11) of lo', praotiedly all of the copper is in solution as the molecular species, and the true solubility product constants are much lower than those in Table 5. From s, pedagogical standpoint, however, no harm is probddy done by this oversimplification.
Copper Salts, Analysis The copper salts may easily be analyzed for waters of hydration by baking a weighed quantity at 1 0 0 T for 3 hr and reweighing. Five of the sdts are monohydrates (I*), the exceptions being the benzoate and the p-anisste which are trihydrates. The latter two compounds undergo interesting color changes on dehydration. Percent water loss and theoretied percentages are listedin Table 6. Copper analysis may be carded out by thermogravimetric analysis of the hydrated salts or, better, by crucible ignition to
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and magnetic susceptibilities of the copper salts may also be measured.
copper(I1) oxide of the anhydrous salts. Results of these procedures are shown in Table 6. The TGA results are errtttic, probably due to small sample size, while the crucible ignitions give uniformly low values for copper, probably due to the vigorous natlire of the decomposition which expels some copper from the crucible unavoidably. In this regard, the nitro compounds are dmost explosive, snd careful ignition is necessary. Far some of the copper salts (0-nitro, p-nitro, o-chloro in psrticular), a dark green crystdine form precipitates from the supernatant on standing. This form, which has identical stoichiometry to the light blue form, has s "binuclear" dimeric structure (IS) as compared to s. polynuclear structure for the normal precipitate. This second crystal form may also be profitably analyzed for percent water and copper. Percent copper may illso be determined by electrodeposition of copper from solutions of the more soluble salts. The low solubility of the para derivatives prevents easy determination of percent copper by this method.
The students who have been involved in this program have demonstrated considerable enthusiasm. Analytical results agree well enough with expected values or are internally consistent enough that self-confidence is generatedin the student. Finally, the burden of a prograni of this type on the laboratory instructor is less than in the case of conventional laboratory programs. There is no necessity for week-to-week mixing of reagents and solutions for the next laboratory period since the students generate all their own chemicals.
Silver Salts
Acknowledgment
The silver salts may be precipitated in s. manner analogous to the copper salts. Again assuming all dissolved silver to be in the form AgC, a solubility product constant may he calculated bv eravimetric means. Values are listed in Table 7. A check
The authors were supported in this work by a grant from the Gustavus Adolphus College Research Fund. Thanks are also due those students who took part in the initial phase of the program.
Conclusions
Literature Cited termined gravimetricdly. Supernatant silver ion concentration was determined by potentiometric measurement also (16)and this result was used to determine mother value for the solubility product constant. These results (Table 7) are very much lower than for either of the two preceding methods, and probably constitute a more realistic value for the solubility product constant, since the measurements are not sensitive to any moleculsr species present. A Leeds & Northrup #I552 potentiometer was used for all determinations, but an ordinary student potentiometer would be quite suitable.
Extension
These exercises may be easily extended to prepara tion and analysis of other salts, Thermogravimetry is easily done if instrumentation is available. Spectra
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Journal of Chemical Education
(1) W ~ ~ s oJ. u .R.. J. Cnmx. Eonc., 46, 447 (1969). (2) N o n n x m . J.. J. C H ~EDOO., . 44, 891 (1967). (3) WALTER R. I., J . CIEM. EDUC.. 45, 673 (1988). (4) ~ o c r a n Ja., : W.c., J . Cnew. E D ~ O . . 45, 523 (1988). (5) ADAMB,R.. J O A N ~ O NJ., R.. A N D WILCOX,JR.. C. F., "Laboratory Erpedmenta in Organic chemistry" (5th ~d.). he ~ a c m i l l a nConpany, New York, 1963, p. 304. (a) H o o a x m . C. D. ( E d r t o ~ )"Tablea . for Identification of Organic Compounds," Chemical Rubber Publiahing Company. Cleveland. 1960, p. 118-134. (7) P ~ A WP,~ I.. , am SIIHINB. R. J. J., Can. I.Chen.. 46,241, (1988). (8) BAD*.J. L..J. Cnmr. Enno.. 46,889 (1969). (9) HBINZ,D. E., J . CHEH.EDDC..44,115 (1987). Eonc., . 41,288 (1884). (10) PmxasroN, R. C . . J . C ~ E M (11) MAY,W. R., AND JONEB,M. M., J. Inow. Nucl. Chern., 24,511 (1986). (12) LEWIS.J., LIN,Y. C., ROYSTON, L. K., AND TAOXPBON, R. C., 3. Cham. Soc., 8484 (1965). T ~R. c., ~~t~~~ ~ 200,468 ~ (1883). ~ ~ (13) L ~ W , ~J., (14) FAZTZ, iatry", J. Ailyn S.. AND& Baaon. S C ~ NInc., RJA.. . Boston, G. H., 1966, Quantitative p. 154. AndYti~sICbem(15) F ~ , WJ., 8..
SOHENK.
JR., G. H., OP. it.. P.448.
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