Complexes of Squaric Acid with Acid-Base Indicators Robert I. GeIb and Lowell M. Schwartz Department of Chemistry, University of Massachusetts-Boston,
Boston, Mass. 021 16
Complexes between the undissociated squaric acid molecule and the unprotonated forms of cresol red and 4-phenylazodiphenylamine indicators have been found. Measurements of the complexation equilibrium constants lead to the conclusion that the indicators have about the same affinity for the undissociated squaric acid as they do for the hydrated proton.
HAVINGMEASURED the primary aqueous dissociation constants of "squaric acid" (1,2-dihydroxycyclobutenedione)by pH potentiometric titration (1) and by conductometric methods (2, 3), we thought it appropriate to confirm these results by one more independent technique, namely by colorimetry. Knowing from our previous work that pK, was about 0.5,we used several readily available indicators: cresol red, thymol blue, methyl violet, ethyl violet, malachite green oxalate, and 4-phenylazodiphenylamine. From measuring the visible spectrum of each indicator as a function of pH in aqueous HC1 solutions, we found that cresol red and 4-phenylazodiphenylamine were suited for our requirements. In both cases, the color change was caused by a simple shift of intensity from one absorption peak to another corresponding to a shift in the indicator equilibrium between protonated and unprotonated forms. Using a standard solution of cresol red, we constructed a pH us. color working curve by measuring the absorbance of the protonated indicator as a function of pH in HCI. Then we measured the visible spectra of several concentrated oxalic and squaric acid solutions containing standard cresol red solution. From pH values read from the working curve and from reasonable estimates of activity coefficients, we calculated a primary dissociation constant for oxalic acid that closely approximated the known value, but a nearly infinite value of K1 for squaric acid. Since experimental uncertainty could not account for this strange result, we hypothesize that the undissociated squaric acid as well as the solvated proton is capable of reacting with the unprotonated form of the indicator and that in the visible region both the protonated indicator and the indicator-squaric acid complex are spectrally similar (Figure 1). However, we observed a small shift of approximately 1-3 nm in the position of the peak near 518 nm. This shift increased with increasing squaric acid concentrations and we ascribed it to a small difference in the wavelengths of maximum absorption of the protonated indicator and of the complex. In effect, we surmise that the undissociated squaric acid appears to the unprotonated indicator essentially as a proton. Method. We analyze the pH behavior of the visible spectra of the indicator cresol red as follows: Cresol red changes color from yellow to orange as the pH drops from 3 to OS. Spectrally this corresponds to an increased absorption centered at 518 nm and a loss of absorption intensity near 435 nm. The 518-nm peak height is taken as a measure of the protonated form of the indicator. (1) L. M. Schwartz and L. 0. Howard, J. Phys. Chem., 74, 4374 (1970). (2) Zbid.,75, 1798 (1971). (3) R. I. Gelb, ANAL.CHEM., 43, lllO(1971). 554
ANALYTICAL CHEMISTRY, VOL. 44, NO. 3 , MARCH 1972
Experiments to be discussed later show that the protonated form is uncharged. The absorbance A of this species is, according to Beer's law A
=
&[HIn]
(1 )
where E is the absorptivity at the maximum, b the absorption path length and [HInl the molar concentration of protonated indicator. The indicator equilibrium is
where aH is the hydrogen ion activity, [In-] the concentration of unprotonated indicator, y- and yo are the activity coefficients of the charged and uncharged indicator species, respectively. If Co is the analytical concentration of indicator (in our work a small but unknown quantity which we estimate at 10-5M),then [In-] [HIn] = Co and Equations 1 and 2 can be combined to yield
+
(3) where CI, ebCo is regarded as an unknown parameter. By measuring the absorbance A of a series of indicator-HC1 solutions, all with the same indicator concentration but various HC1 concentrations, and plotting 1/A us. yo/aay-, an absorbance vs. pH working curve will be obtained which, if it is a straight line, will verify that Beer's law applies to the HIn absorbance and that proper activity coefficients y o and y- were assumed. The slope of this line will be KI,/CI, and the intercept on the ordinate axis will be I/CI~ which values determine both KI, and CI,. The effectiveness of the cresol red indicator in this application may be tested by measuring the absorbance in oxalic acid solutions containing the same (CO) concentration of indicator. By referring to the working curve, the value of yo/aay- is found and then by an iterative procedure the ionic strength, degree of dissociation, and the primary dissociation constant of oxalic acid is calculated and compared with the known value. According to our hypothesis, the absorbance of indicatorsquaric acid solutions at 51 8 nm will be A = sb[HInj
+ ecb[HzSq.In-l
(4)
where [HzSq.In-] is the concentration of squaric acid-indicator complex and e, is the absorptivity of this species. If the complexation reaction is written as an association in the form HzSq In- t3 H2Sq.In-, then the equilibrium constant is
+
where yc- is the activity coefficient of the complex. Since yc- and y- both refer to similarly charged rather large molecules their values were taken to be equal and Equation 5 was used in the simpler form which appears at the right. Because the total indicator concentration is very small, the primary dissociation remains
Figure 1. Absorption spectrum of cresol red solutions A , 0.2000 MHCI; B, 0.1000MHCI; C, 0.0920M squaric acid; D, 0.1277M oxalic acid; and E , 0.0030MHC1
600 WAVELENGTH
IN
NU
correction term. We see that Equation 1 should properly be written
A Here we have assumed that a mean activity coefficient y + suffices for the singly-charged hydrogen and bisquarate ions and that the secondary dissociation (pK2 = 3.5) ( I ) has negligible effect on the essential equality of [H+] and [HSq-1. Csa is the analytical squaric acid concentration. In these solutions the total indicator is distributed three ways and so C" = [In-] [HIn] [HzSq.In-]. Combining this with Equations 2,4, and 5 , we obtain
+
+
where Cc ecbC". If we wish to use this relationship as a means of characterizing the squaric acid-indicator complex, it is necessary to determine KI, and CI,from Equation 3 using data obtained from a series of absorption measurements with the indicator at known acid (HCI) concentrations; to measure A in a series of indicator solutions of varying squaric acid concentration ; and to plot G us. A and extract a slope of -Kc and a G intercept of CcKcaccording to Equation 7. Other values needed for Equation 7 are [HzSql and aH which by knowing Kl for squaric acid are calculable from Equation 6 and the various activity coefficients which must be estimated from semiempirical correlations. If, as we suggested earlier, undissociated squaric acid interacts with the unprotonated indicator in the same manner as does a hydrogen ion, then we would expect the HzSq .In- association constant Kc to be approximately the reciprocal of Kin, the H+In- dissociation constant, and the parameter C,( = ecbCo)to be approximately equal to CI,(= cbC"). Ancillary Considerations. A number of problems are associated with the analysis of the hypothesized complexation equilibrium as described above. First, the spectrum of cresol red as a function of pH (see Figure 1) shows that at 518 nm, the center of the absorption peak due to the protonated form, there is also extraneous absorption from the tail of the peak corresponding to the unprotonated form. The Beer's law relationship written as Equations 1 and 4 does not explicitly account for this and must be modified by a
= eb[HInl
+ db[In-I
(8)
where e ' is the absorptivity due to the tail of the In- peak at the center of the HIn peak. The concentration of In- may be eliminated from this equation by using the expression C" = [HIn] [In-l, the result being
+
A,,,,
= A - E ' K " = (E - e')b[HIn]
(9)
The difference (e - E') is a constant and may be regarded as an effective absorptivity. Consequently, it may be replaced by another symbol, ciz., e. The term e'bC" represents a correction to A . The numerical value of this term is found experimentally by measuring the absorbance of the indicator at a pH sufficiently high so that [HIn] is negligible. For cresol red in pH 4 solution, the peak due to the protonated form has negligible amplitude, and the absorbance at 518 nm is due only to the unprotonated form, its value being e'bCO. This is independent of pH and is simply subtracted from all absorbance values read at 518 nm to yield corrected absorbances A,,,,. Equation 4 can be similarly treated by adding the absorbance e'b[In-] so that A
=
eb[HIn]
+ e,b[HSq.In-] + t'b[In-]
(10)
The concentration of In- may be eliminated from this equation by utilizing the expression Co = [HIn] [HZSq. In-] [In-l, and
+
A,,,, = A - e'bC" =
(e
- e')b[HIn] (ec
+
+ - e')b[HzSq.In-l (11)
Again (e - e') and (ec - e') may be regarded as effective absorptivities and replace e and ec in Equation 4. A,,,, then effectively replaces A in all preceding equations. The evaluation of activity coefficients is done with semiempirical correlations. Activity coefficients of the smaller ions such as hydrogen ion, binoxalate ion, and bisquarate ion are estimated utilizing the Davies equation (4)
(4) C. W. Davies, "Ion Association," Butterworths, London, 1962. ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972
555
~~~
~
~
~~
Table I. Comparison of pH Values Measured Potentiometrically with a Glass Electrode and Colorimetrically with Cresol Red
Solution Oxalic acid Squaric acid
Molarity 0.07619 0.1008 0.1277 0.07544 0.1086 0.1632
Cresol red Cresol red from from pH meter Equation 3 Equation 7 1.37 1.28 1.23 1.27 1.14 1.00
1.389 1,307 1.235 1.216 1.083 0.951
...
...
:
1 276 1,145 ,998
where D = 0.51 at 25 ’. This equation has been shown by Davies to represent y values of fifty 1:l electrolytes up to 0.1M to within 1.6%. For the larger ions, the charged indicator and complex species, we used the Debye-Huckel (5) equation
where B = 0.33 A-1 and the “ion-size” constant ai is regarded as an empirical parameter. The value of ut that is eventually found is consistent with reasonable estimates of molecular dimensions. The activity coefficients 70of the uncharged species are assumed to vary as (6) log 70= f k Z (14) where k is an empirical parameter whose value is assumed to be 0.1 for undissociated squaric and oxalic acids but is determined experimentally for the uncharged indicator species. Before the activity coefficient parameters can be determined, it is necessary to establish the charges on the protonated and unprotonated forms of the indicator. The species charges written into Equations 1 and 2 , Le., [In- and HIn] were chosen from the other likely alternatives, [In and HIn+] and [In+ and HIn2+]by measuring the absorbances of a series of 0.1M HC1 in indicator solutions with KCl added to vary the ionic strength up to 0.24M. (The possibility [In2and HIn-] was rejected because the cresol red molecule has only two possibly acidic protons and exhibits another color transition near pH 8 so that the presence of In2- species at low pH values seemed unlikely.) Equation 2 can be rearranged as
E .-. -K I J H I ~ I Yo
aH[In-l
and the other two alternatives, if written in comparable form, would be
y+- - K I , [ H I ~ ~ + I y2+
uH[In+l
KI, is independent of ionic strength, UH decreases with ionic strength in a known manner, and [HInl/[In] = A/(ebCo - A ) varies accordingly through A . In the series of experiments, the right-hand side of Equations 2a, b, and c decreased with (5) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,”
second ed. (revised), Butterworths, London, 1965. (6) J. N. Butler, “Ionic Equilibrium,” Addison Wesley, Reading, Mass., 1964; W. F. McDevit and F. A. Long, J . Amer. Chem. Soc., 74, 1773 (1952).
556
ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972
increasing ionic strength. This behavior is consistent only with the possibility y-/yo considering that y + / ~ and + ~ yo/y+ are both increasing functions of ionic strength in this range, but y - 1 ~ 0is a decreasing function of the ionic strength. The parameters ai and k in Equations 13 and 14 are interdependent but can be evaluated by two series of experiments and an iterative calculation. The first series is the measurement of A us. [HCII. An estimated value of k is chosen to calculate yo and then Equation 3 plots are made with various at values. The ut which yields the best straight line is selected along with the corresponding KI, and CI,. The second series is A us. [KC11 mentioned in the previous paragraph. Data from these measurements are plotted in the form log 7 0us. I where
(% log yo
=
log
1) aHy-
KIn
(3a)
and the slope k is evaluated. This new value of k is used to refine the calculation of ai, KI,, and CI, which are in turn used to refine k . We carried out these calculations with the aid of digital computer programs. To explore whether the hypothesized squaric acid complexation is unique with cresol red or a more general phenomenon, we carried out a series of experiments and calculations using 4-phenylazodiphenylamine (PDPA) as the indicator. The limited solubility of this substance in water was overcome by adding enough acetone to bring C” up to a photometrically convenient level. Consequently, all experiments with PDPA were done in 17.8 wt % acetone solution and the proper activity coefficient equations had to be reconsidered. Since the Davies correlation Equation 12 was developed entirely for aqueous solutions, we used the DebyeHuckel form given by Equation 13 exclusively. Since the parameters D and B in this equation vary with solvent dielectric constant 6 according to D = 6-3/2 and B a 6-’12, the dielectric constant of 68.3 for this solution ( 7 ) leads to values of D = 0.63 and B = 0.35 A-l. The ion-size parameter ai was taken as 6 A for the smaller ions and 12 A for the indicator and complex ions. Ionic strength experiments revealed that the PDPA indicator species charges are In and HIn+ so that the activity coefficient ratio (y-/yo) in Equations 2, 3, and 7 were replaced by (yo/y+). Finally the calculation of K, from Equations 6 and 7 requires a knowledge of K , the primary dissociation constant of squaric acid which is unknown in 17.8% acetone. We measured this value as 0.667 + 0.004 and the corresponding constant for oxalic acid as 1.714 =t 0.006 by a conductometric method (3). In order to rule out the possibility of some gross error in solution preparation or in our interpretation of the colorimetric data, we measured the pH values of both oxalic and squaric acid solutions potentiometrically with a glass electrode which had been standardized against our 0.1M HC1 solution and against a buffer with pH = 4.00. The standardizations agreed within the estimated readability of the meter. Measured pH values of the oxalic and squaric acid solutions were consistent with the known pK values for these acids within experimental uncertainty, 0.02 pH unit. The potentiometric values are compared with those obtained from photometric data in Table I. Two methods were used to calculate the pH values in squaric acid solutions. The first involved direct use of Equation 3 to obtain UH from A and in effect ignored the H2Sq.In- complex. The second (7) G. Akerlof, ibid., 54, 4125 (1932).
/
/
Table 11. Corrected Absorbance Values (Cresol red at 518 nm, 4-phenylazodiphenylamineat 537 nm) Series 1. Cresol red (correction constant c’bC” = 0.019) Molarity &,* HC1
0.300 0.2000 0.1Ooo 0,0650 0,0300 0,1000 0.1OOO
Oxalic acid
Squaric acid
0.407 0.365 0.291 0.237 0.153 0.286 0.282
+ 0.0304M KC1 + 0.1433M KCI
0.2034 0,09430 0.08738 0.04590
0.291 0.224 0.217 0.162
0.1561 0.1219 0.08080
0.341 0.314 0.271
Series 2. Cresol red (correction constant ebC”
HC1
Figure 2. Plot of Series 2 0 and Series 4 HCI-indicator data in the form 1/A plotted as the ordinate us. yo/aEy- for Series 2 or us. y+/aEyofor Series 4 according to Equation 3. The lines are the calculated least-squares best fits
Oxalic acid
Squaric acid method accounted for the formation of the complex through Equation 7. Most of the calculations were done by digital computer. The program to calculate the Equation 3 plot fitted a straight line to the 1/A us. y&aE data by the method of leastsquares. Its output was the least-squares line slope and intercept, the RMS deviation of the ordinate values from the line, and the parameters K I , and CI,. Another program calculated the dissociation constant of oxalic acid from those two constants and a series of absorbance us. oxalic acid concentration data. A third program accepted the two constants, the absorbance us. squaric acid data, and the primary dissociation constant of squaric acid and from these calculated a least-squares straight line of the data in the form of Equation 7 and gave K , and C,. In this program the activity of hydrogen ion was found in each solution by an iterative calculation in a sequence involving initial estimation of I , then of y- from Equation 13, y o from Equation 14, aH from Equation 3, and y& from Equations 12 or 13. These values in turn, gave a refined estimate of I = a&* which was employed in the next iteration.
0.039)
0.751 0.687 0.594 0.541 0.429 0.318
0.1277 0 . 1008 0,07619
0.519 0.476 0.427
0.1632 0.1359 0.1086 0.09200 0.07544
0.681 0.654 0.608 0.571 0.531
Series 3. Cresol red (correction constant E’bC” HCI
=
0.040)
0.1365 0,09100
0,684 0.584
0.1436 0.1247 0.1020
0.690 0.665 0.618
Squaric acid
Series 4. PDPA (correction constant c‘bC” HC1
=
0.014)
0.1820 0.1365 0.09100 0.07280 0.04550 0.09100 0.039MKCl 0.09100 0.078MKC1 4 0.09100 + 0 . 1 5 6 ~KC1
0.756 0.660 0.507 0.439 0.315 0.519 0.520 0.535
0.1675 0.1397 0.1139 0.1071 0.08222 0.07248 0.06832
0.422 0.382 0.338 0.334 0.279 0.266 0.247
0.1207 0.1053 0.08994 0.07458 0.05923
0.613 0.568 0.517 0.460 0.389
+ +
Oxalic acid
EXPERIMENTAL
Absorbance measurements were made with a Cary 14 recording spectrophotometer thermostated at 25 + 1 “C and employing borosilicate glass IO-!nm cells with conductivity water as the reference. (Closer temperature control was unnecessary since no temperature variation of our data could be detected in this range.) Squaric acid was purchased from Aldrich Chemical Company and dried under vacuum at about 100 “Cfor an hour. The acid samples were titrated with standardized NaOH and found to have an equivalent weight within 0 . 6 z of the theoretical value of 57.03. Cresol red was purchased from Fisher Scientific Company and other indicators were obtained from Eastman Chemical Company. These were used without further purification. Aqueous solutions of hydrochloric, squaric, and oxalic acids (together with whatever impurities were present in the reagents) were
=
0.2000 0.1500 0.1oOo 0,0800 0,0500 0,0300
Squaric acid
~~~~
each found to be completely transparent to visible light so that extraneous absorbance from these solutes did not interfere with the indicators’ visible spectra. Table I1 shows the results of the absorbance measurements on four complete series of experiments each with a different indicator concentration. The tabulated absorbance values are corrected according to Equation 9 or 1 1 . ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972
557
Table 111. Summary of Calculated Results Oxalic acid, K In CI, KC CC pK, i std dev Series 1 (cresol red) 0.05629
0.5345
31.84
0.3986
1.226 i 0.007
31.24
0.7469
1.214 i 0.004
29.78
0.8224
...
10.03
1.171
1.59 j z 0.03
Series 2 (cresol red) 0.05403 1.079 Series 3 (cresol red) 0.05647 1,126 Series 4 (PDPA) 0,1439 1.315
Quantity A y
(Eq. 12 or 13)
Table IV. Error Analysis Estimated Propagated uncertainties in Kr n CIn uncertainty 0.005 1.5% nil 1% nil 1s
a , (Eq. 13)
2A
0.4%
0.9%
Maximum estimated uncertainty KI, CIn
pK, squaric acid (Eq. 12 or 13) HzSq impurity
y
3% 1% 0.03 1% 0.6%
Maximum estimated uncertainty
7.0
K,
CC
15% 15% 3% 6% 2.4%
0.2% 1.5% 0.7%
42%
4%
nil 1,2%
Oxalic acid pKl SERIES
60
KIn CI,,
3% 1%
A
0.005
50
y (Eq. 12 or 13)
1%
2
0.04 0.03 0.04 0.01 0.12
7,: 7.0
4.0
6.5-
\T
S E R I E S ‘1
6.0-
1 \
Figure 3. Plot of Series 2 0 and Series 4 0 squaric acidindicator data in the form G plotted as the ordinate us. A according to Equation 7. The lines are the calculated leastsquares best fits RESULTS
The full visible spectra of a number of Series 2 solutions are reproduced in Figure 1 to show how the intensity shifts between the peaks as the p H changes and that the squaric acidindicator spectrum has essentially the same form as the others but is shifted slightly to higher wavelength. The least-squares straight lines through the HCl-standard solution data plotted according to Equation 3 were found relatively insensitive to changes in the activity coefficient parameter at but for Series 1 the line with the sma!lest RMS deviation was calculated when at was set at 12 A , and this value was used in all subsequent series for indicator species ut. The calculation of the parameter k in Equation 14 from ionic strength experiments yielded essentially k = 0 for cresol red and k = 0.4 for PDPA. The results of Equation 3 and Equation 7 calculations are given in Table 111. The data points and the calculated least-squares straight lines for the HCl standards and the squaric acid solutions are shown in Figures 2 and 3, respectively, for both Series 2 and 4. 558
We note that the three series using cresol red are in agreement with each other in predicting the indicator dissociation constant KI, and the squaric acid complex association constant K,. These two constants and the corresponding values from Series 4 are not quite reciprocals as would be the case if the indicator had equal affinities for the hydrogen ion and the squaric acid molecule. The affinities of both cresol red and PDPA appear slightly greater for the squaric acid than the hydrogen ion although this difference may be accountable by the experimental uncertainties in the values. The two parameters CI, and C, are not exactly equal for each series as would be true if the absorptivity of the complexes and the protonated indicators were identical. However, the ratio CC/Crn( = E ~ / E I , ) is nearly constant for the three series of cresol red experiments, its value being 0.72 f 0.03. The indicator (cresol red) determination of p K for oxalic acid in seven aqueous solutions yielded values averaging 1.22 and with a standard deviation of 0.02 which should be compared with the value 1.27 measured by Darken (8) conductometrically. The comparable oxalic acid data in 1 7 . 8 x acetone are pKl = 1.59 i 0.03 measured colormetrically in Series 4 and pK, = 1.71 + 0.01 measured conductometrically (3) in this laboratory.
ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972
DISCUSSION
With digital computer programs available for most of the calculations, it was a simple matter to find the effect of experimental measurement uncertainties on the computed parameters by rerunning the program with slightly perturbed input data. Table IV lists our estimates of the principal uncertainties inherent in these calculations and the resultant uncertainty propagated to the calculated parameters. We estimate that the experimental uncertainty in each A datum due to photometric inaccuracy and reading error amounts to 0.005. The corresponding uncertainties in the coordinate values in Figures 2 and 3 are shown as error flags drawn on the individual points. These account for the random scatter of the points around the straight lines drawn in (8) L. S . Darken, J. Amer. Chem. SOC.,63,1007 (1941).
feasible. If as this study apparently demonstrates, squaric acid forms complexes with acid-base indicators which are colorimetrically nearly the same as the protonated indicators, other strong undissociated acids perhaps do likewise. Since oxalic acid and squaric acid species are structurally similar, it would not be surprising to find an oxalic acid-indicator complex. If such a complex were formed with a small association equilibrium constant, a colorometric determination would yield abnormally high values of hydrogen ion concentration and correspondingly low pKl values.
these figures. All the other entries in Table IV represent systematic errors. Within our total estimated uncertainty of 42z in K,, the equilibrium constants K, and K I , are reciprocals. Apparently the squaric acid-indicator complex and the hydrogen ion indicator complex are equally strong. Our error estimates, however, do not account for the excess of CI,over C,in each series and so we must conclude that the protonated indicator has a slightly greater absorptivity than the squaric acid complex. We notice that the pKl values for oxalic acid determined colorimetrically here are consistently lower than the values determined conductometrically. The discrepancies might be accountable by the 0.12 estimated uncertainty in the colorimetric values. However, another explanation seems
RECEIVED for review July 16, 1971. Accepted November 4, 1971. Acknowledgement is made to the donors of the Petroleum Research Foundation administered by the American Chemical Society for partial support of this work.
Rapid Method for Quality Appraisal of Various Compounds by Quartz Crystal Thermometry Flora C . Youngken Central Research Department, Experimental Station, E.I. du Pont de Nemours and Co., Wilrnington, Del. 19898 The potential of quartz crystal thermometry for general laboratory application to quality appraisal by means of freezing point or melting point is discussed. Attention is drawn to the practicality of assigning a dual laboratory function to the referenced cryoscopic 8, instrumentation and method. Studies of benzene and dimethyl sulfoxide substantiate the degree of differentiation which can be ascertained among specimens of these materials without procedural tedium or involved calculations. The reported data illustrate the outstanding features of a quartz crystal thermometer with respect to reproducibility, sensitivity, and high “C) of temperature readings. resolution (1 X Values of absolute accuracy are established by an easy, uncomplicated calibration procedure. A comparison of literature freezing point values with those determined for six standard aqueous NaCl solutions with a concentration range of 0 5 5 % demonstrate the mean error in our results. Design of a tube and stirring rod suitable for determinations using 5 ml of sample is included.
THE IDENTICAL EQUIPMENT assembled for the cryoscopic determination of number-average molecular weight serves a second analytical function in our laboratory. With this apparatus, quality appraisals by means of freezing point (FP) or melting point (MP) determinations are accomplished by the same technique that we use for establishing a FP value for M n determinations ( I ) . A digital indicating quartz crystal thermometer (QCT) as the temperature-measuring device was the prime instrumental innovation in our previous report. On the basis of an investigation, in depth, of the solvents, benzene and dimethyl sulfoxide (DMSO), for application to M n cryoscopy, augmented by our experiences with quality studies via FP or MP of other materials, we now recommend the use of a QCT for this new purpose in general laboratory operation.
(mn)
(1) J. S . Fok, J. W. Robson, and F. C. Youngken, ANAL.CHEM., 43, 38 (1971).
The subject of cryoscopy and FP or MP as a purity criterion is well covered by Glasgow and Ross (2). The possibility of the detection of minute F P differences among specimens of the same product is dependent upon the sensitivity and precision of the temperature-measuring tool available to the analyst. Furthermore, if temperature values of absolute accuracy are necessary, a means of certification of these data is required. Research Paper 1676 (3) from the National Bureau of Standards (NBS) describes a procedure for the determination of the purity of hydrocarbons by measurement of freezing points. Since the publication of this report in 1945 and up to the present time, this particular time-temperature apparatus has become one of the most widely used in this country ( 4 ) . In 1963, Ross and Dixon proposed changes in technique and apparatus to improve the original cryometric method of purity determination for very highly purified materials, but the Glasgow-Streiff-Rossini techniques were the only ones which gave excellent agreement for all samples of benzene to which controlled amounts of contaminants had been added for a cooperative investigation by 20 participating laboratories in 1961 ( 5 ) . The thermometric system described in Research Paper 1676 consists of a 25-ohm platinum resistance thermometer, a Mueller-type resistance bridge with main coils thermostated, and a highly sensitive galvanometer (1 mm on the scale adjusted to an equivalence of 1 to 5 X 10-4 “C). Soon after publication of this procedure, it became the basis of a standard method of the American Society for Testing and Materials (ASTM). Essentially the same temperature-measuring set(2) I. M. Kolthoff et al., “Treatise on Analytical Chemistry,” Interscience, New York, N.Y., 1968, Part 1, Vol. 8, Chap. 88. (3) A. R. Glasgow, Jr., A. J. Streiff, and F. D. Rossini,J . Res. Naf. Bur. Stand., 35 (6), 355 (1945). (4) Ref. 2, p 5053. (5) G. S. Ross and H. D. Dixon, J . Res. Nat. Bur. S t a ~ d 67A, , 247 (1963).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972
559
