for comparison. These sensitivity values, in agreement with a precise definition of sensitivity in trace analysis (IO), are 12 X 108 mm.M-' and 3 X loe mm.M-' for the titration with catalytic end-point indication [malachite green-periodate ionmanganese(I1) system] and for the indirect kinetic method, and for volumes of 10 and 15 ml, respectively. If a similar treatment is performed with titrations of EDTA with Cu(I1) as catalytic titrant and 1-ascorbic acid-Oz indicator reaction (11, I2), S(lo-~) is equal to 5.5 X lo6mm.M-'. This is half the sensitivity of the malachite green-periodate-Mn(I1) system. It should be mentioned that even though the titra(10) J. E. Barney 11, Tufuntu,14,1363 (1967). (1 1) H. A. Mottola, M. S. Haro, and H. Freiser, ANAL.CHEM., 40, 1263 (1968). (12) H. A. Mottola, K. Muller, and H. Freiser, unpublished results, 1967.
tion with Cu(I1) can be performed at a pH between 6 and 7, the catalyst must be used at rather high concentrations to get well developed curves. This works against the low limit of detection expected by using a rather high pH. ACKNOWLEDGMENT
The author gratefully acknowledges the assistance of Steven E. Henley in some of the experimental work.
RECEIVED for review December 22, 1969. Accepted February 24, 1970. Paper presented at the 17th Anachem Conference, Detroit, Mich., September 1969. This work was supported by the Oklahoma State University Research Foundation and the National Science Foundation (Grant GP13472).
Calorimetric Determination of EquiIibrium Constants for Very Stable Metal-Ligand Complexes Delbert J. Eatough' Shell Development Company, Emeryuille, Cal$ It has been previously demonstrated that titration calorimetry can be used to obtain reliable equilibrium constants for complex metal-ligand systems if the K values for the stepwise addition of the Ii ands are less than approximately lo4. A procedure or the calori. metric determination of Kvalues for very stable metalligand systems i s given in this work. Log K, AH", and A S O values have been calorimetrically determined for the interaction of Hgz+ion with 2-aminoethanol and of Cu2+and Zn* ions with 1,lO-phenanthroline in aqueous solutions at 25 O C . For these systems the equilibrium constants for the formation of the various species vary from 106 to lolo. The calorimetric results are in good agreement with previously reported values determined by potentiometric techniques and demonstrate that titration calorimetry can be used to quantitatively determine the thermodynamics of interaction for any metal-ligand system if the proper choice of titrant can be found.
P
less than about lo4. Equilibrium constants for several complex systems where the various K values are less than l o 4 have been reported (2, 4), and excellent agreement is found between the calorimetrically determined values and values determined by more conventional techniques such as potentiometry or spectrophotometry. However, no calorimetrically determined equilibrium constants have been reported for any metal-ligand system where K is greater than lo4. Equilibrium constants greater than lo4 could be reliably determined if a competitive equilibrium of the type,
is found such that KR values for the stepwise removal of the ligand, L, from MLn by N are in the region I KR I lo4. If the equilibrium constant for the reaction N+L=NL,
CALORIMETRIC determination of equilibrium constants requires that significant concentrations of both reactants and products be in equilibrium during the titration of one reactant with another (1-3). If the equilibrium constant for the interaction is large enough that the products are stoichiometrically formed as the reactants are mixed, then the thermometric titration curve will be linear and independent of the magnitude of the equilibrium constant for the interaction. This restriction limits these systems which can be studied by titration of a metal species with a ligand to those for which the K values for stepwise addition of the ligand to the metal are Present address, Center for Thermochemical Studies, Brigham Young University, Provo, Utah 84601 (1) J. J. Christensen, R. M. Izatt, L. D. Hansen, and J. A. Partridge, J . Phys. Chem., 70,2003 (1966). (2) R. M. Izatt, D. Eatough, R. L. Snow, and J. J. Christensen, ibid.,72, 1208 (1968). (3) J. J. Christensen, D. P. Wrathall, and R. M. Izatt, ANAL. CHEM., 40, 175 (1968).
KN
is independently known, then the K values for the formation of the ML, species from M and L could be obtained by this method. The application of this technique to the determination of p K values in the region 4-9 for a single protonation step has been demonstrated (3). If the principle can be reliably applied to the study of complex systems, then the possible systems for which equilibrium constants may be determined calorimetrically is greatly increased. Using titration calorimetry techniques we have determined stepwise equilibrium constants, enthalpy and entropy change values for the interaction of Hg2+ion with 2-aminoethanol, A , to form HgA2+ and HgA22+ and of Cu2+and Zn2+ ions with 1,lO-phenanthroline, P, to form MPZf, MPz2f, and MPZ2+. The competing ion, N, in all cases is the hydrogen ion. These systems were chosen for the study for the following reasons: (1) Accurate p K values valid at 25 "C and zero ionic strength, (4) D. J. Eatough, Ph.D. Thesis, Brigham Young University, Provo, Utah, 1967. ANALYTICAL CHEMISTRY, VOL. 42. NO. 6, MAY 1970
635
p,
have been reported for the protonation of A (5-9) and P
(IO-M), (2) Reliable equilibrium constant values have been reported for each of the metal ligand systems (15-24) with which the results of the present study can be compared, (3) Because all the interactions involve the addition of a neutral ligand to the metal ion or hydrogen ion, the equilibrium constants andpK values are expected to be independent of p, thus simplifying the calculations, and, (4), ThepK value for ionization of the HA+ ion is greater than the log K values for the stepwise interaction of Hg2+ion and A while the pK value for proton ionization from the H P ion is smaller than the log K values for the stepwise addition of P t o either Cu2+ or Zn2+. Thus by this choice of ligands, systems for which the KR values are both greater than and less than unity can be studied. A calorimetrically determined A H " value for protonation of A at 25 "C has been reported (25), but no A H " values have been reported for the interaction cf A with Hg2+. Andregg (26) has reported calorimetrically determined A H " values at 20 "C for the interaction of H+ and Cu2+ with P but no A H o values valid at 25 "C have been reported for the various phenanthroline systems. J?,XPERIMENTAL
Materials. Reagent grade HgO (Baker analyzed 100.0% red powder), ZnO (Mallinckrodt Analytical, lOO.Ox),CuNQ3. 3 H z 0 (Baker Analyzed, 99.7 Z), HCIOl (Baker Analyzed), H N 0 3 (Mallinckrodt Analytical), 1,lO-phenanthroline HzO (Eastman White Label), and 2-aminoethanol (Eastman White Label) were used in the preparation of the solutions. The metal salts were assumed t o be as analyzed. The 2-aminoethanol was further purified by distillation before use. The NaOH solutions were standardized against acid potassium phthalate (National Bureau of Standards Acidimetric Standard). The HCIOa and HNOa solutions were standardized against the NaOH solutions and the 1,lO-phenanthroline and 2-aminoethanol against the acid solutions. Sufficient HClOa o r H N 0 3 was added to the Hg(C104)2r Cu(N0J2, and Zn(NO& solutions t o suppress hydrolysis of the various metal ions. Freshly boiled distilled water was used in e
(5) R. G. Bates and G. D. Pinching, J . Res. Natl. Bur. Stand., 46 349 (1952). (6) J . R. Lotz, B. P. Block, and W. C. Fernelius, J . Phys. Chem., 63,541 (1959). (7) J. Belin and G. Douheret. Comot. Rend.. 261. 984 (1965). (8j S. P. Datta and A. K. Grzybowski, J . Chern.' Soc., 1959, 1091. (9) G. Douheret and J. C. Pariad, J. Chim. Phys., 59, 1013 (1962). (10) R. Nasanen and E. U. Usitalo, Suomen Kem., 2 9 , l l (1956). (11) H. H. Perkampus and H. Kohler, 2. Electrochem., 64, 365 (1960). (12) S. C. Lahari and S. Aditya, 2.Physik. Chem., 41,173 (1964). (13) A. A. Schilt, J . Amer. Chem. Soc., 79,5421 (1957). (14) C. H. Cook, Jr., and F. A. Long, ibid., 73,4119 (1951). (15) J. Bjerrum, Chem. Reu., 46,381 (1950). (16) J. Lotz, unpublished results, 1959. (17) C. V. Banks and R. I. Bystroff, J. Amer. Chem. SOC., 81, 6153 (1959). (18) G . Andregg, Helu. Chim. Acta, 46,2397 (1963). (19) J. R. Scharff and M. R. Paris, Compt. Rend., 263, 935 (1966). (20) R. J. Pflaum and W. W. Brandt, J . Amer. Chem. SOC.,76,6215 (1954). (21) J. M. Dale and C. V. Banks, Znorg. Chem., 2, 591 (1963). (22) I. M. Kolthoff, D. L. Leussing, and T. S. Lee, J . Amer. Chem. SOC.,73,390 (1951). (23) H. Irving and D. H. Mellor, J. Chem. SOC.,1955,3457. (24) M. Yasuda, K. Sone, and K. Yamasaki, J. Phys. Chem., 60, 1667 (1956). (25) D. L. Levi. W. S. McEwan, and J. H. Wolfenden, J . Chem. ' SOC.,1949, 766 (26) G. Andregg, Helu. Chim. Acta, 46, 2813 (1963). 636
ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
the preparation of all solutions and all solutions were stored and used under a pure nitrogen atmosphere. Procedure. All runs were made in a Tronac Inc. constant temperature environment titration calorimeter at 25.00 "C. The method of data analysis is similar to that previously described ( I ) . The A H " values for protonation of A and P were first determined by titration of solutions of A and P with HCIOa. A precipitate was formed during the titration of P so this value was checked by titration of a NaOH solution with a PsHNOa solution. The A H " value obtained from the latter titrations was 0.4 kcal/mole less exothermic than that obtained by titration of P with HC104, presumably due to the interaction of Clod- ion with the HP+ ion. The A H " value obtained from the titration of NaOH with P."01 was taken to be the correct value. Nitrate was used as the anion in the C U ~and + Zn2+titrations t o avoid the interaction between HP+ and Clod- ions. The interaction between Hg2+ and A was studied by titration of two different HgA2(C104)z-HC104 solutions with HCIOa. Two different Cu(NO& and Zn(NO3)2 solutions were each titrated with a P-HN03 solution. Corrections for dilution of the HClOa titrant were calculated from literature data (27) and dilution corrections for the P. H N 0 3 titrant were determined by titration into N a N 0 3 solutions having the same p value as the metal nitrate solutions. Calculations. The procedures used in analysis of the data ( I ) and the least squares calculation of equilibrium constant and enthalpy change values from thermometric titration data ( 2 ) have been previously described. The AH" values for protonation of the A and P species were calculated using the pK values 9.498 (5) and 4.857 (IO), respectively. Corrections for the formation of HzO in the various runs were made using literature values for the ion product (28) and heat of ionization (29) of HzO. It was necessary t o make corrections for hydrolysis of a small amount of Hg2+ ion in the studies involving the interaction of Hg2+ and A . These corrections were made using the values log K = -3.22 and A H = 6.6 kcal/mole for Reaction 3. Hg2+
+ H60 = H@H+ + H+
(3)
The HgOH+ species is the only appreciable hydrolysis product under the conditions of the present study (30). The K value used to correct for hydrolysis of the Hgz+ ion was calculated using the value log K = 0.36 at p = 0 (31,32) for the reaction Hg(OH)z Hg2+ = 2HgOH+ and the value log K = -6.80 for Reaction 4 valid at p = 0. This latter value was calcu-
+
+
+
Hg2+ 2Hz0 = Hg(0H)z 2H+ (4) lated from literature values for the hydrolysis of HgClz (33) and for the interaction of Hg2+with C1- (34, 35). The A H o value was calculated from reported calorimetric data valid at p = 3.0 (36) for Reactions 3 and 4 and p = 0 (33, 37) for (27) C. E. Vanderzee and J. A. Swanson, J . Phys. Chem., 67, 285 (1963). (28) H. S. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," 3rd ed., Reinhold Publishing Corporation, New York, N. Y., 1958. (29) J. D. Hale, R. M. Izatt, and J. J. Christensen, J . Phys. Chem., 67, 2605 (1963). (30) I. Ahlberg, Acta Chem. Scand., 16, 887 (1962). (31) A. B. Garrett and W. W. Howell, J. Amer. Chem. SOC.,61, 1730 (1939). (32) K. Damm and A. Weiss, 2.Naturforsch., 106, 534 (1955). (33) J. A. Partridge, R. M. Izatt, and J. J. Christensen, J . Chem. SOC.,1965, 42311 (34) L. G. Sillen and A. E. Martell, "Stability Constants of Metal' Ion Complexes," Special Publication No. 17, The Chem. SOC., London, 1964. (35) L. Hansen, R. M. Izatt, and J. J. Christensen, Znorg. Chem., 2, 1243 (1963). (36) R. Arnek and W. Kakolowicz, Acta Chem. Scand., 21, 1449 (1967). (37) R. J. P. Williams,J. Phys. Chem., 58, 121 (1954).
Table I.
Log K, AH', and A S o Values Valid at 25 "C in Aqueous Solution Comparison with Previous Literature Values
H+
+
Hg2+
Reaction A = HA+
+ A = HgA2+ +A
H+ + P
=
=
HgAz2+
HP+
+ P = CuP2f
CUP2+
AH' (kcal/mole)
-11.54
=t
0.03
ASo (gibbslmole)
4.7
0.1
+P
CUP22+
=
CUP22+
+ P = CUP32+
8.56 f 0.07 (8.51)'66 (8. 99)16 8.77 i 0.05 (8.81)'54 (8.71)18 (4.857)'O
9.14 It 0.06 (9.08)17 (9.12)'BC (9,16) (6. 30)zo 6.87 =t 0.08 (6. 68)17 (6,68)1*c (6.96) (6. 15)L0 (6. 42y1 5.42 =t 0 . 1 (5.18)" (5.14)'BC
(-12.1)' (- 11. 5)6 -10.2 f 0 . 5
+ P = ZnP2+
ZnP2+
0 ? 0
5.0 It 2
-8.3 zk 0.5
12.3 f 2
-3.60 f 0.02
10.1 It 0.1
+ P = ZnP22+
(4.63)21 6.17 It 0 . 1 (6.83)21 (6. 46)lBC (6.36)" (6.43)*% (6. 5)23 5.91 It 0 . 1 (5.22)Z' (5.74)180 (5.64)'' (5.72)22
0.4 1 .o 0 0.4 1.0 0 0
(-3.5)" (-4.1)'s (- 3.95)*6* -11.03 i 0.1
0.4 0 0.1 4.8 f 0.3
0
0.01 0.1 0.5
0.4 -5.42 zk 0 . 1
13.2 f 0 . 3
0
0.01 0.1 0.5
0.4 0.1 -5.1 i: 0.3
1.1 i 0.9
0
0.01 0.1 0.4 -7.5 f 0 . 2
3.1 f 0.6
0.1 0
0.1 0.1 0.1 0.1 0.1 -4.76 f 0 . 1
11.1 =t 0.4
0 0.1
0.1 0.1 0.1
0.1 0.1
(5.4923
(5.9124 5.25 f 0 . 1 (4.87)s' (5.16)'s~ (5.20)17 (4.85)ZZ (5.10)23 (4. 8)24
0 0
(- 3 * 3 ' 0
(5.50)
Zn2+
P
(- 12.07)"
HgA2+
CuZ+
Log K (9.498)s
-2.9 =t 0.7
14.3 it 2
0 0.1
0.1 0.1 0.1 0.1 0.1
The log K values have been converted from 30 "C using the AH' values determined in the present study. Calorimetrically determined at 20 "C. c The log K values have been converted from 20 "C using the AHo values determined in the present study.
Reaction 4. The hydrolysis correction is from 1 t o 2 % of the total measured heat and the calculated log K values for interaction between Hg2+and A are changed by less than the standard deviation by changes of 0.5 log units in the K value for Reaction 3. A Debye-Huckle expression of the form,
where the various terms have their usual significance (28), was used to convert equilibrium constants valid at p = 0 t o concentration quotients were necessary. The values a o =
4A" and C = 0.2 were used in the calculations (38). Calculations were aided by the use of a UNIVAC 1108 computer. RESULTS AND DISCUSSION
The results of the present study, together with previously determined literature values are given in Table I. I n all cases the expressed uncertainty is the standard deviation between runs. The thermometric titration data from which the results given in Table I were calculated have been deposited with the (38) G. H. Nancollas, "Interactions in Electrolyte Solutions," Elsevier, New York, N. Y., 1966. ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
637
Time, minulei
Figure 1. Titration of a C u ( N 0 & - H N 0 3solution with P.HNO3 National Auxiliary Publications Service (NAPS) of the American Society for Information Science (ASIS) as document number 00935. A microfiche ($2.00) or photoprint ($5.00) copy may be obtained from NAPS c/o CCM Information Sciences, Inc., 22 W. 34 St., New York, N . Y. 10001. The agreement between the present results and previously determined equilibrium constants is in general very good, showing that titration calorimetry techniques can be used to accurately determine equilibrium constants for the formation of very stable metal-ligand complexes. The calorimetric values fall well within the range of previously reported values, with the exception of the Zn2+-P system for which the K, value is about 0.2 log unit lower and the K2 value is about 0.2 log unit higher than most previous results. The K3 value for this system is, however, in good agreement with previously reported values. The calculated log Kl and log K2 values for the interaction of ZnZ+ ion with P agree well for the two Zn(NOs)z solutions used, indicating there are no serious systematic errors in the present study ( 3 , 39, 40). Considering the wide range of previously reported values and the complexity of the system the agreement is satisfactory. The log K value for the formation of Cupz+ from Cu2+and P is 4.3 log units larger than the p K value for proton ionization from H P . Theoretical considerations on systems where a single equilibrium is occurring suggest that a log K difference this large could not be determined t o better than 0.2 log unit under the conditions of the present study (3, 39, 41, 42). The precision of the log K1 value given in Table I for the interaction of Cu2+and P is ~ t 0 . 0 6log unit and the agreement of Kl with previously reported values is excellent. If only one equilibrium of the form given by Equation 1 occurs during the titration the products will be essentially quantitatively formed (39) J. J. Christensen, D. P. Wrathall, J. 0. Oscarson, and R. M. Izatt, ANAL.CHEM., 40, 1713 (1968). (40) J. J. Christensen, J. H. Rytting, and R. M. Izatt, J. Chem. SOC. ( A ) , 1969,861. (41) D. P. Wrathall, Ph.D. Thesis, Brigham Young University, Provo, Utah, 1967. (42) S. Cabani and P. Gianni, J. Chem. SOC.( A ) , 1968, 547. 638
ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
as titrant is added t o the calorimeter when log KR is greater than about 3.5. I t is thus expected that log KR values larger than 3.5 cannot be precisely determined using titration calorimetry. If however successive equilibria occur during the titration the formation of the various products may not be stoichiometric and equilibrium constants larger than log KR = 3.5 may be precisely determined. A correlation between the heat produced due t o the interaction between Cu2+ and HP+ during the titration and the corresponding species present in the calorimeter is given in Figure 1 for the titration of a Cu(NO& solution with P.H N 0 3 . As Figure 1 indicates, significant concentrations of both Cup2+ and CuPZ2+are present at the point when the Cu2+:P ratio in the calorimeter is 1:l. Because the AH" values differ for formation of the two species, this situation results in significant curvature in the titration curve in this region making possible the calculation of both Kl and Kz from the calorimetric data. It then appears that the range over which K values may be determined may be larger for systems with simultaneous equilibria than for a system involving only a single equilibrium. The log K and AH" values for the various metal ions given in Table I were calculated assuming the p K and AH" values for proton ionization from the protonated ligand are correct. Any error in these two values will be directly reflected in the calculated log K and AH" values for the corresponding metal ion complexes because the concentration of free ligand in the solution is always negligible. Thus, for example, if the p K value for HA+ is 10.00 rather than 9.50, the corresponding log K value for the interaction of Hg2+ ion with A to form HgA2+ would be calculated to be 8.56 0.50 = 9.16. The competing ion in the present study is the hydrogen ion; however, the study could have been carried out using another metal ion to study the equilibria. F o r example, the log K values for the interaction of Ag+ ion with P to form AgPf and AgP2+ are known (34) and slightly larger than those for the interaction of Zn2+ion with P, so that titration of a ZnP32+ solution with Ag+ would yield the equilibrium constants for the interaction of Zn2+with P. The equilibrium constants obtained by such a titration would be expected to be less precise than these given in Table I because of the greater complexity of the system. IfKvalues for theAg+-Pinteractions were not known, relative values for reactions of the type given by Equation 1 could still be determined by such a titration. I n a similar manner, the proper combination of titrant and titrate solutions would yield relative K and A H values for the interaction of a ligand with any two metal ions, protonation of any two ligands, or the interaction of a given Lewis acid with a pair of Lewis bases. The simple systems such as the one studied in this report or previously reported by Christensen et al. ( 3 ) would be expected to give the better values. Because the ability to study a chemical system is limited only by the requirement that a significant amount of heat be produced during the titration, titration calorimetry should prove to be a valuable tool for the quantitative study of metal-ligand or acid-base interactions in either aqueous or nonaqueous solvent systems. Comparison of the present results with the log K and AH' values for the interaction of C U ~(+2 ) and Zn2+ ( 4 3 ) with pyridine, Py, provide a n excellent system for studying the bonding of a chelate. The chelate effect has been defined by Schwarzenbach (44) as the added stability conferred upon a complex containing bidentate o r higher order ligands as
+
(43) D. J. Eatough, R. M. Izatt, and J. J. Christensen, Brigham
Young University, unpublished results, 1967. (44) G. Schwarzenbach, Heh. Chim. Acta, 35, 2344 (1952).
Table 11. Unitary Chelate Effect for 1,lO-Phenanthroline and Bipyridine at 25 “C Reaction AGO (kcal/mole) A H o (kcal/mole) -.TASo (kcal/mole) CuPy2+ P = cuP2+ 2Py -4.2 -2.2 0.4 CuPy2+ 2P = CuP,a+ 4Py -8.9 4.9 -12.8 znPy,2+ P = znP2+ 2Py -0.9 -7.4 6.5 znPyaa+ 2P = znP2=+ 4Py -4.4 -5.7 1.3 CUPY,’+ B = CUB’+ 2Py -1.2 0.5 -1.7 CUPY,’+ 28 = CuBzz+ 4Py -4.3 4.7 -8.0 ZnPyZz+ B = ZnB2+ 2Py 1.6 3.5 -2.0 ZnPya*+ 2B = ZnB22f 4Py 0.0 10.9 -9.8
+ + + + + + + +
+ + + + + + + +
compared with a complex containing the corresponding unidentate ligand. It has been pointed out (45)that comparison of thermodynamic data to study structural effects is valid only for the mole fraction concentration scale. The thermodynamic quantities associated with the chelate effect for the interaction of Cu2+and Zn2+ with P and P y are given in Table 11. Values for the chelate effect of bipyridine B, are also included where the log K and A H o values used for the interaction of Cu2+ and Zn2+ with B are those reported by Atkinson and Bauman (46)valid in lrn NaC104. The thermodynamic quantities given in Table I1 have been converted to the mole fraction scale by the method previously outlined (45). The results given in Table I1 indicate that in the case of bipyridine the chelate effect is entirely due t o the favorable entropy changes, the A H o values being endothermic in all cases. For the reaction of Zn2+ with B the A H o values are (45) J. J. Christensen and R. M. Izatt in “Physical Methods in Advanced Inorganic Chemistry,” H. A. 0. Hill and P. Day, Eds., Interscience, New York, N. Y., 1968. (46) G. Atkinson and J. E. Bauman, Jr., Inorg. Chem., 1, 900 (1962).
endothermic enough that the chelate stabilization disappears when the thermodynamic quantities are put on the mole fraction basis. For the 1,IO-phenanthroline interactions the trend is largely reversed and the chelate effect is mainly due t o favorable enthalpy changes. The stabilization of chelates due to favorable enthalpy changes has been previously noted for amine chelates (47). It is not presently clear why the bipyridine system apparently differs from the 1,IO-phenanthroline and other nitrogen containing chelate systems. For both ligands the chelate stabilization is significantly larger for the Cu2+ complexes than for the corresponding Zn2+ complexes. ACKNOWLEDGMENT
The author thanks Donald C. Guthrie for assistance in performing the thermometric titrations. RECEIVED for review November 24, 1969. Accepted February 24, 1970. Appreciation is expressed to Shell Development Company for release of the material for publication. (47) A. E. Martell, Advan. Chem. Ser., 62,272(1967).
Improved Instrumentation for Phosphorimetry of Organic Molecules in Rigid Media Ruth Zweidingerl and J. D. Winefordner Department of Chemistry, Unirjersity of Florida, Gainesville, Fla. 32601 Phosphorimetry has previously been of limited use as a result of marginal precision and accuracy as well as difficulties and time of sampling. As a result of a rotating sample cell, a more stable source power supply, and a better solvent clean-up procedure, detection limits have been lowered more than one hundred-fold, precision and accuracy have been increased by more than ten-fold and a considerable reduction in time and effort in the sampling and measurement procedure has resulted as compared to phosphorimetric measurements with standard commercial equipment. The excellent precision and detection limits have been obtained not only in clear rigid solvents but also in cracked glasses and snowed matrices. The possibility of performing precise, accurate, sensitive, selective, rapid analysis in solventsformingopaque or densely cracked glasses greatly extends the usefulness of phosphorimetry-particularly to samples of biological interest.
PHOSPHORESCENCE IS A TYPE of photoluminescence in which radiation is emitted by an organic molecule following excitation of the molecule by ultraviolet or visible radiation of higher Research Triangle Institute, Research Triangle, North Carolina.
energy than the emitted radiation. It is experimentally differentiated from fluorescence, the other major type of photoluminescence, by its longer lifetime and lower energy. Phosphorescence is a true molecular property and so is useful in the identification of organic molecules. The mechanism of the production of phosphorescence is well-known ( I , 2) and will not be discussed here. Phosphorimetry is the analytical method involving the use of phosphorescence for quantitative analysis. Zander (3) and Winefordner, McCarthy and St. John (4) have described the use of phosphorimetry in analysis. Phosphorimetry has been used during the past decade for a limited number of analyses, e.g., the analysis of impurities in ~
~
~
~~~
~
(1) D.M. Hercules, “Fluorescence and Phosphorescence Analysis,” D. M. Hercules, Ed., Interscience, New York, 1966. (2) C. A. Parker, “Photoluminescence of Solutions,” Elsevier, New York, 1968. (3) M. Zander, “The Application of Phosphorescence to the Analysis of Organic Compounds,” Academic Press, New York, 1968. (4) J. D. Winefordner, W. J. McCarthy, and P. A. St. John, “Methods of Biochemical Analysis,” D. Glick, Ed., Vol. 15, Interscience, New York, 1967. ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
639
