Standards Analytical Instrumentation Applications of instruments in the past may seem fairly trivial in today's world of pulsed and Fourier transform spectroscopy. B u t at the time, even these contributions were much appreciated. Dr. Kamm, sometime director of research at Parke, Davis and expert in qualitative organic, once told us how a chemist had worked 11 years on the structure of cholesterol and received the Nobel Prize for it before it was found to be the wrong structure. With present equipment, we could probably get a pretty good handle on t h a t compound in a few hours. I had the good fortune to be injected into analytical instrumentation quite early. T h e new "mechanical" techniques were sometimes viewed skeptically. Acceptance of them had to be earned, whence came the preoccupation with instrument performance and standards. Standards were significant then and still are. The term standards applies to all our operations. There are standards at the bench, of the instrument, of its performance, of the data it produces, a n d so on. Many of the examples to follow will be taken from spectrophotometry because it was a bellwether of the time, and my major interest. When the photoelectric spectrophotometer first came on the scene, there was great enthusiasm on the p a r t of those of us who had done absorption by earlier methods. Spectra could be obtained so easily and quickly! But frustration began to develop as spectra were collected on many types of samples because relationships with chemical structures were obscure. It would be hard to believe now the excitement generated when Woodward showed t h a t cv,/i-unsaturated ketones
have characteristic high intensity absorption with wavelengths affected predictably by substitution. It is probable t h a t this neat little correlation got him more instant attention, more invitations to speak at meetings, and so on, than did any of his later honors. T h e drug laboratories soon caught on to the quantitative potential for spectrophotometric measurement of isolates and purity. It was virtually magic to be able to check capsule fill simply by dissolving the contents a n d reading concentration by absorption. Remember t h a t at this time the criteria available and accepted for qualification of a sample were the melting point, the ash content determined by burning a spot of it on a spoon, and sometimes the specific rotation. This splendid new assay approach was not well accepted by the drug authorities. For one thing, it was not a legally certified procedure. Also, work in their laboratories and elsewhere demonstrated a lamentable lack of consistency. T h e inference was t h a t maybe the method was not dependable. Their reaction was to propose t h a t for each drug assay, a sample of certified standard be run concurrently by the same procedure. T h e n if the spectrophotometric result for the standard differed from a previously determined value for t h a t material (how this would be established was not specified), the assay value was to be corrected by a corresponding amount. This doubled the work, and in a sense the uncertainty. It was unfavorable too t h a t some drug standards were difficult to keep in quality state, as a result of hydration, instability, or whatever. It seemed to us t h a t it would be a
386 A · ANALYTICAL CHEMISTRY, VOL. 49, NO. 4, APRIL 1977
better approach to show t h a t the spectrophotometer was operating effectively. One method for this was checking it by means of absorbance readings on solutions of primary absorption standards, quantitatively prepared. Potassium chromate in alkaline solution was a favorite. Potassium dichrom a t e in acidic solution has a better spectrum, with a number of maxima and minima in the absorbance range. But neither is stable over long periods, a n d both are so intensely absorptive t h a t it is difficult to prepare appropriate solutions without a dilution step. Before long, the recording spectrophotometer came into use, and now the problem was even more acute. T h e automatic nature of its operation tends to separate the operator from an intuitive appreciation of its performance. So long as the lights come on and the wheels turn, it is used. T h e problem was specific. We needed an absorption standard which could be read directly in the instrument to give assurance of its performance quality. T h a t standard would best be solid, of optical quality, stable to light and handling, easily cleansed, economic, and capable of being duplicated in quantity in all shapes and sizes for different species of spectrophotometers. Further, its use must require no change whatever in instrument adjustments or settings, require no auxiliary equipment, and above all require no interpretation or calculation of the data obtained. And it must be quick and easy, or it would not be done. Did any such paragon exist? Why surely. After working through perforated screens, mesh screens, and so on, appropriate glass filters were clearly
Report
J. M. Vandenbelt Division of Research and Medical Affairs Parke, Davis & Co. 2800 Plymouth Road Ann Arbor, Mich. 48106
the answer. Having arrived at this, we learned that Gibson and coworkers at the National Bureau of Standards had made a careful and thorough study and arrived at the same conclusion. A difficulty was their coverage of the subject in such minute detail that the impact of their contribution was lost in the mass. There is a lesson for us here: In our reports and announcements, let us hold to a Standard of Succinctness, lest the peaks of value be lost in clouds of verbiage. Now, how necessary is this standardization? Are spectrophotometers really so shaky that we need to spend time and effort just to see how they run? In one of the early tests, we sent a set of filters to 10 drug laboratories with recording spectrophotometers, all of which were fairly new, all operated by trained, experienced, and concerned specialists. Of those 10 instruments, five were found to be performing well, three were defective in some particular (including our own), and two were badly off norm. Subsequent studies have given similar results. Instruments do require periodic check with suitable standards. At Parke, Davis, we instituted a regular run on all the recording spectrophotometers at five locations. If an instrument slumps, it is detected soon, and correction made. The next twist in instrument development was the recording spectrophotometer with near-infrared capacity. We needed a standard for this region also. A little careful study made it evident that the rare-earth didymium filter, lauded in World War II for Winnie-the-Welder's eye protection, would be just fine if cut to about four
times standard thickness. I requested a quotation for this from Corning and received a two-page letter of reasons why this would not be possible. Shortly thereafter, we happened to be visiting with my sister, who then lived in Corning. I drew as partner a fiend at bridge who turned out to be special products manager at Corning Glass. I told him my story and in a couple of weeks received a beautiful specimen. This was sent around to collaborating laboratories; it was held up unreasonably at one address. It developed that they had two infrared instruments, one of which responded normally to the peaks and valleys of the filter spectrum. The other gave a bland, almost zero tracing. That second instrument was being used to monitor an isolation program of aliphatic lipids, and its complete failure to show appropriate bands had wasted months of work. That laboratory was most responsive to collaborative studies after that. We still needed a filter standard for the ultraviolet. Learning that Corning had developed a new glass transmitting well in the ultraviolet, I asked an acquaintance in research there to melt some potassium chromate into it. This he did, working extra at night, and obtained a piece of attractive jade green. Components of the melt had reduced the orange chromium VI to green chromium III, which was of no value in the ultraviolet. Therefore, he tried it again in an oxidizing environment, with splendid results. Unfortunately, this filter did not meet Coming's fantastically high standards of light stability. We tried to no avail to show that the tiny exposures we would give it in the dispersed beam of a double
monochromator for a few minutes would not amount to a millionth of their test requirements. We still do not have a good filter standard for the ultraviolet. Some of the field service engineers for spectrophotometers carry a small tube of benzene vapor to check the wavelength adjustment after repairs. We looked for a more substantial alternate, and it seemed that the rare earths have the richest endowment of narrow spectral lines. Of all of these, holmium has the best array in both ultraviolet and visible. My friend at Corning was induced to melt holmium oxide in that glass, as a minimal hope of obtaining an absorbance standard with peaks and valleys in the ultraviolet. Everybody knows that an absorption spectrum in a constricted environment shows reduced resolution. The vapor spectrum of a compound has more resolution than its solution, and the liquid state more than the solid. Contrary to expectations, the bands of holmium were not at all broadened in the glass, but yielded a spectrum with spikes of absorption so sharp that their precision, it is said, exceeded the wavelength devices of the time. The Bureau of Standards took it over, and you can get from them a certified holmium wavelength absorption filter (Figure 1). Operating standards are helpful, but they do not go far enough. There must be also a Standard of Knowledge of the properties, responses, and ranges of every instrument in use, and operation must not exceed those limits. Some years ago a paper originating in Australia reported that control measurements of biphenyl required a
ANALYTICAL CHEMISTRY, VOL. 49, NO. 4, APRIL 1977 · 387 A
Figure 1. Absorption spectrum of holmium in glass
concentration calibration. In other words, it did not obey Beer's law. In addition to maligning one of my favorite compounds, biphenyl should be about the last in the world not to obey Beer's law. It has no substituent groups; there is no way for solvation, ionization, or other effects to occur. We compared readings of biphenyl at increasing concentrations on an in strument like the author's and found t h a t biphenyl did obey the law within the proper range of the instrument. We looked also at higher concentra tions with an instrument of larger range, with the same result. T h e au thor had failed to note the optimum range of his instrument. Once at a national meeting we lis tened to a paper by a well-known spectroscopist. T h e room got very quiet as it became evident, to some at least, t h a t the theoretical basis of his discussion was based on false low wavelength maxima due primarily to scattered light. T h e compounds them selves could have no absorption maxi ma in this region. T h e chairman must have realized this too, because at the end of the paper he called immediate ly for the next one, giving no chance for questions or discussion. It is in cumbent on every responsible opera tor to know the characteristics of his instrument and to see t h a t its ranges are not exceeded. Other types of instruments can ben efit from standards too. A few years ago we were monitoring a program in volving a product whose most signifi cant characteristic appeared to be the specific rotation. Fortunately, we had just acquired one of the better new spectropolarimeters. But if a prep failed to show the anticipated quality, you know what happened. "What's
the matter with that instrument this t i m e ? " "How come it doesn't check yesterday?" In self-defense we had to devise a suitable operating standard. T h e traditional procedure for check of polarimeters is to read a quantita tive solution of a primary standard compound. T h e old, old, literature, in fact, specifies this as a necessary oper ation for any new polarimeter: Pre pare a series of graded concentrations and calibrate the instrument scale from the readings. But a solution standard will not do for long-term consistency check or for transport to other participating units. We looked for a better way. The handbooks list a great many compounds with rotation, most of which are solid, rare, expensive, and unavailable. There are a few substitut ed aliphatic liquid esters of potential interest. Examinations of stockroom samples of these were disappointing, in that they were mostly racemic, probably having been prepared by synthesis. But a few of the old dusty bottles in the back shelves were more interesting; their contents were ob tained long ago from natural sources. One of these, 2-methyl-l-butanol, gave a fine response. Readings on this standard were remarkably consistent,
giving a fine check for the polarimeter, and most convenient for comparison readings with the distant instruments. An unexpected bonus is that the com pound has a very low-temperature coefficient of rotation. It appears to be completely stable, and it is available on stock order for a few dollars from Columbia Organic Chemicals Co., Inc. Since having it, we can check the po larimeter in a minute. It would be out of place here to be labor Standards of Vigilance in bench work, with which we are all acquaint ed. However, it is documented that some very casual spectrophotometric values are published. Professor Phil lips compared the published absorp tion data for a number of simple or ganic compounds such as benzoic acid, and found a variation u p to 25% in just one volume of his "Organic Elec tronic Spectral Data". We can only hope t h a t many of these were reported by workers or their assistants in disci plines other than analytical chemistry. Beyond knowing the instrument and its operation, there is another and limitless area we can call Standards of Understanding. Here is the largest op portunity of all to make a real contri bution, both for your laboratory and for yourself. To do this, search out
Figure 2. Formula for calculation of isoelectric point Source: J. P. Greenstein and M. Winitz. "Chemistry of the Amino Acids", Vol Ι, ρ 483, Wiley, New York, NY.,
388 A · ANALYTICAL CHEMISTRY, VOL. 49, NO. 4, APRIL 1977
Figure 3. Graph of percent group ionization with pH
methods of Clarification and Simplification, find the Kernels of Principle, and Pursue Correlations. It is probable t h a t the better examples of Clarification and Simplification are forced on us by the complexity of traditional methods. Here's an example. We were involved in a peptide synthesis program in which it was desirable to know the isoelectric point p i ' of the products. T h e formula is calcu-
lated simply, as shown in Figure 2, provided you are confident of getting the values in the right slots. Every ionization takes place over a practical range of four p H units, the center point of which is the pKa of the group (Figure 3). R a t h e r than struggle with the complex formula, we can plot the ionizations of glycine, for example, showing charge behavior of the groups vs. p H (Figure 4). T h e carboxyl group is neutral at low p H but acquires a
Figure 4. Group ionizations of glycine, glutamic acid, lysine 390 A · ANALYTICAL CHEMISTRY, VOL. 49, NO. 4, APRIL 1977
negative charge as the p H increases to about 4. Meanwhile, the amino group remains fully protonated u p to p H about 8, where it begins to lose the cation. T h e isoelectric range (1ER) of the compound is over the whole range p H 4 - 8 . T h e published isoelectric point, midway between the two p H values, is a most inadequate designation of the true picture. When there are two groups of the same type vs. one of the other, as in glutamic acid, the carboxyls acquire charges at different pH's. At t h a t p H where the sum of both negative ionizations is unity, t h a t only is the p H at which positive and negative charge on the molecule is equal. T h e result is self-evident by inspection, and there is unequivocally a specific isoelectric "point". Similarly for lysine, the p H a t which the sum of the positive partial charge on the two amine groups is unity just equals the negative charge on the carboxyl, giving again a specific isoelectric point. With tyrosine (Figure 5), the isoelectric range is decided by the carboxyl and amine ionizations, because the hydroxyl group on the phenyl cannot get into the act until higher p H . B u t if the phenol group is substituted with iodine, its p K a is reduced so t h a t a negative charge begins to form at about p H 6. T h e isoelectric range is limited to that region before the hydroxylate begins to form. With diiodotyrosine the pKa is reduced further, resulting in a small isoelectric range. A simple representation like this can be
Figure 5. Group ionizations of tyrosine and derivatives
Figure 6. Absorption spectra of benzene and derivatives with acceptor groups
most illuminating and convincing. Turning now to Standards in Un derstanding, we should touch on some of the correlations of electronic spec tra and chemical structure by Leonard Doub and myself. Standards of Hon esty and Scientific Accuracy require t h a t he be given lead billing in ben zene correlations. T h e opportunity to work with him on these projects was one of the fortunate accidents of a ca reer. T h e following will touch on some principles of the subject. These studies include an analysis of the characteristics and behavior of the benzene chromophore as affected by substitution, group type, group ioniza tion, and steric combinations. T h e spectrum of benzene is shifted to longer wavelength by substitution as a function of the donor or acceptor strength of the substituent group. T h e primary band shift with typical groups of increasing acceptor strength is shown in Figure 6. Note t h a t this is not the resolved low intensity finestructure band usually associated with benzene, but the intense band with some resolution at about 200 nm. The effect of carboxyl group ioniza tion is shown for benzoic acid. Ioniza tion of the carboxyl adds a negative charge, which reduces the attracting power of the group. As a result, the primary band shifts to lower wave length. Similarly, the spectrum of aniline is drastically changed on formation of the amine cation in acid. In this state, its electrons are immobilized with re spect to the ring, and the spectrum re verts to t h a t of unsubstituted benzene
Figure 7. Absorption spectra of 4'-hydroxyacetophenone and its anion
itself. This occurs with all amine de rivatives and is often a most valuable clue to identity and a basis for analy sis. Note t h a t the benzene primary is shifted to longer wavelength by sub stitution with groups of either type. If groups of opposing types are present in the 1,4-position, their combined ef fect is substantially larger than the sum of their individual effects. By ob serving the wavelength displacement of each group when interacting with a group of the opposite type, we have determined the individual displacing effects of the common groups in t h a t environment. T h e spectrum of 4'-hydroxyacetophenone is a typical example of this interaction (Figure 7). T h e spectrum is considerably displaced from t h a t of acetophenone, whose maximum is about 240 nm, due to the interaction across the chromophore with the hy droxy group. If the hydroxyl is ion ized, it becomes a very strong donor, so there is a pronounced wavelength shift. There is present also a new sec ond primary band at lower wave length, brought out from the lower ul traviolet by the shift. Numerical derived displacing values were worked out for the common sub stituent groups as shown for 4'-aminoacetophenone and ρ-anisic acid in the formula δλό Χ δ\"0 = 24.05 (λ 180) (Figure 8). Predicted vs. observed wavelengths for a large number of compounds are generally less than 5 nm in error. These were determined first in aqueous solvent because the state of
392 A · ANALYTICAL CHEMISTRY, VOL. 49, NO. 4, APRIL 1977
group ionization was best known and controlled in this medium. Later the correlation was developed for other solvents, especially methanol. As a re sult of the fortunate circumstance t h a t displacement of the two group types is affected oppositely in solvents other
Figure 8. Calculation of derived displac ing values in 4'-aminoacetophenone and 4'-methoxybenzoic acid
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Figure 9. Absorption spectra of 2'- and 3'-hydroxyacetophenones and anions
Table I. Derived Displacing Values of Substituent Groups (δλ0 in nm) for Two Solvents Type I (donor)
Type II (acceptor)
H20
MeOH
OCOCH, CH3 F CI Br OH O© OCH3
26 29 24.5 29.3 31.5 36 55 37
26.5 29 25 29.5 32 38 59 38
NH 4 ® S02H S02CH3 S02NH2 AsO(OH)2 CN CONH2 CONHNH2
°^0>
37
37.5
CONHNH3®
NH2 NHNH2 NHCH3 NHC2H5 N(CH3)2 N(C 2 H S ) 2 NHCONH
50 40 57 58 61 63 45
54 42 59.5 61 62 64 47
C02h
Ν
0
56
58
Ν
ΝΗ
R
59
Ô
63
•α
66
R
co 2 © C02CH3 COCH, CHO N02 NO
H20
MeOH
25 33 38 38 38 44 46
25 32.5 37 37 38 43 46 44.5 49
50 44 50 63 69 90
47 42 48 59 63 81 92
Figure 10. Absorption spectra of antibiotic A-65 and 4-nitrotoluene
than water, the same formula can be used with other solvents also. Derived displacing values for the common groups are tabulated in Table I. Combinations of the stronger groups shift the primary nearly to the visible region. Familiarity with these sequences provides a framework and fabric of orientation, around which one can organize one's thinking and spectrum files very like an alphabetical index. If the opposing groups are in the 1,2- or 1,3-positions, an entirely different type of spectrum results (Figure 9). T h e bands increase in intensity from long to shorter wavelength. T h e secondary band is clearly evident and
Figure 1 1 . Absorption spectra of reduced antibiotic, 4-methylaniline, and amine cations
shifted to comparatively long wavelength. The first primary band shifts moderately as an additive function of the group displacements. T h e second primary, which is substantially less intense than the first primary with 1,4disubstitution, here is intense indeed. It is most unexpected on various counts t h a t these isomer pairs would have such similar spectra. Similar displacements take place in polysubstituted compounds. If there is a strong 1,4-interaction of donor and acceptor groups, an appropriate band will be present. 1,2- and 1,3-Interactions are indicated by a pronounced long wavelength secondary. These correlations are of daily prac-
Figure 12. Absorption spectra of " S u n d a r e " and analogous structures
tical use in evaluation of ultraviolet absorption spectra. One of the very earliest applications involved a new antibiotic. T h e spectrum by inspection was indicative of a 4-nitrotoluene derivative as characterized by the broad form of the longer band and its wavelength (Figure 10). The inflection at low wavelength is consistent with 1,4-disubstitution. This assignment was received with great reservation, not only because there was no previous record of a nitro derivative from natural sources, but also because the nitro group is not favored in therapeutics. However, if there were a nitro group as indicated, it should be easily reduced with zinc and hydrochloric acid. T h e reduction product would then show a spectrum of different quality very like that of para-toluidine. As a further check, the toluidine spectrum would be shifted markedly to lower wavelength on formation of its amine cation. The derivative did these things exactly (Figure 11). In a few hours, we had established a substantial part of this new antibiotic. A few days later we received for examination an antibiotic from an academic laboratory. Alarmingly, it was the same one. But on learning that we did know a great deal of its structure, that laboratory decided not to work on it and relinquished all priority to Parke, Davis. T h a t antibiotic was chloramphenicol, and our experience put spectral applications in a good position with the company. A last application should be of interest. Many years ago following our correlation work, we were challenged to come up with a sunscreen agent with requisite qualities of light absorption, stability, and water insolubility, but it could not be an amine. All the simple para-disubstituted compounds with primary band intense absorption in
ANALYTICAL CHEMISTRY, VOL. 49, NO. 4, APRIL 1977 · 397 A
the erythemal range require a strong donor group vs. a strong acceptor, and only the amines are donor enough. We did not find a candidate. Then a February or so ago, we ar rived in Florida and were taken direct ly to the golf course. T h e temperature was 95 and the sun a blazing orb, a combination for frightful sunburn. My host delved into his bag and brought out something labeled "Sundare". It was spectacularly effective. In several days of direct exposure, not even a pink developed except for small missed areas. Look at what this inventor did in designing " S u n d a r e " (Figure 12). He
added a vinyl group to the resonance chain. This group adds about 40 nm to the displacement, which allows the use of the comparatively weak but stable and innocuous methoxy group on one end, and the strongly attracting carboxy group on the other. T h e combina tion brings it directly into the erythe mal range. T h e ester group is added for water insolubility and compatibil ity with vehicle substrate. Bibliography J. M. Vandenbelt, J. Opt. Soc. Am., 44, 641 (1954). J. M. Vandenbelt and C. H. Spurlock, ibid., 45,967 (1955).
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J. M. Vandenbelt, ibid., 50, 24 (1960). J. M. Vandenbelt, ibid., 51, 802 (1961). J. M. Vandenbelt, ibid., 52, 284 (1962). J. M. Vandenbelt and Carola Henrich, An alyst, 79,586(1954). K. S. Gibson, Nat. Bur. Stand. Circ, 484 (1949). J. M. Vandenbelt, Analysentechnische Berichte, 36E (1974),Bodenseewerk Perkin-Elmer & Co. GmbH. J. P. Phillips, Anal. Chem., 34, 171 (1962). J. P. Greenstein and M. Winitz, "Chemis try of the Amino Acids", Vol I, p 483, Wiley, New York, N.Y., 1961. A. Albert and E. P. Serjeant, "Ionization Constants of Acids and Bases", Methuen, London, England; Wiley, New York, N.Y., 1962. Leonard Doub and J. M. Vandenbelt, J. Am. Chem. Soc, 69, 2714 (1947). Leonard Doub and J. M. Vandenbelt, ibid., 71,2414 (1949). Leonard Doub and J. M. Vandenbelt, ibid., 77,4535(1955). Anachem Award Address presented at the Third Federation of Analytical Chemistry and Spec troscopy Societies Meeting, Philadelphia, Pa., November 15-19, 1976.
J o h n M. Vandenbelt is senior re search scientist in the Division of Re search and Medical Affairs of Parke, Davis & Co., Ann Arbor. He has been head of its physical chemistry group since 1942. Dr. Vandenbelt received his AB degree in 1934 from Hope Col lege, his AM in 1936 from Boston Uni versity, and his P h D in 1940 from Michigan State University. He is the author of some 50 research papers on physical measurements of compounds, including the properties and behavior of chromophores, spectrum-structure correlations, benzenoid spectra, group ionization and analysis, spectrophotometric standards, the holmium filter, dissociation constants, and protein binding. Dr. Vandenbelt is a member of the American Chemical Society, the Instrument Society of America, the Optical Society of America, the Asso ciation of Analytical Chemists, and the AAAS (Fellow, 1963).
