not interfere and even sharpened the color change slightly when present in amounts of betmeen 10 and 50% of the total solution volume. However, the presence of any base such as ammonia or amines, Rhich might be used to decompose the Grignard complex formed in the preparation of the ketimines, interfered with the analysis. A new method of dccornposing the Grignard complex reported by Tolbert (fa) avoids the presence of the ammonia or amines. The Grignard complexes themselves do not interfere. The presence of more than 1% of nitrile in the titration mixture produces erratic results. Nitriles are usually separated from the ketimines by distillation of the mixture at reduced pressures on a spinning-band column. Ketimines have been titrated in acetic acid in these laboratories to determine purity and yield and to study hydrogen-
ation rates (2, 11, 1 2 ) . The following have also been studied in this investigation or in conjuction with the above researches: diethyl ketimine, w-cyclohexylpentyl 2-butyl ketimine, 2-butyl o-tolyl ketimine, and phenyl fenchyl ketimine. ACKNOWLEDGMENT
The authors thank T. L. Tolbert, ’. H‘ Jenkins, Jr.J D* J * and Young for the samples of ketimines and corroborative
E*F*Englesi
c. w*
information. LITERATURE CITED
(1) Conant, J. B., Werner, T. H., J . Anz. Chenz. Soc. 52,4436 (1930). (2) Dulaney, C. L., Ph.D. thesis, University of Oklahoma, 1956. (3) hIoureau, c., lfignonac, c., -4rzn. chim. 14,322 (1920). (4) Siederl. J., Niederl, V., “Organic
Quantitative
Microanalysis,”
lViley,
( 5New ) Pickard, York,p.1948. L., ANAL, CHEJf. 21, 1015 (1949). (6) Pickard, P. L., Engles, E. F., J r , J. Am. Chem. SOC.75, 21.18 (1953). ( 7 ) Pickard~p. L., Jenkins, s. H., Ibid., 75,5899 (1053). (8) Pickard, P. L., Vaughan, D. J., Ibid., 72,876 (1950). (9) Riddick, J. A , , ANAL.CHEW 24, 41 (1952). (10) Seaman, \IT., -4llen, E., Ibi&23,592 (1951) (11) Soinenfeld, R. J., Ph.D. thesis, University of Oklahoma, 1987. (!2) Tolbert, T. L., N.S.thesis, University of Oklahoma, 1956. Jr.j
RECEIVED for review November 20, 1958. Accepted March 13, 3959. Analytical and Inorganic Section, Second Tetrasectional Meeting, ACS, Ponca City, Okla., March 3, 1956. From a thesis presented by Frank A. Iddings in partial fulfillment of the requirements for the degree of master of rcirnce at the University of Oklahoma, June 1956.
Volumetric Determination of Calcium in Presence of Phosphate RICHARD G. YALMAN and WILLIAM BRUEGEMANN Chemisfry Deparfmenf, Antioch College, Yellow Springs, Ohio PAUL T. BAKER Deportment o f Biophysics, Pennsylvania Sfafe Universify, University Park, Pa. STANLEY M. GARN Fels Institute, Yellow Springs, Ohio
b An indirect titration method using (ethylenedinitri1o)tetraacetic acid is described for the determination of calcium in the presence of phosphate. Small amounts of magnesium and iron d o not interfere. Good results are obtained in the determination of calcium in teeth, bone, and phosphate rock. Less satisfactory results are obtained in the analysis of urine.
T
for the determination of calcium in the presence of phosphate is based upon the precipitation of calcium oxalate (II), which is either titrated with permanganate or converted to the carbonate or oxide. Improvements in the procedure have centered about the mechanics of centrifuging and washing the precipitate (4, IS, 22). A recent chemical development has bren to dissolve calcium carbonate in boric acid and back titrate with standard hydrochloric acid (19). At best the procedures are long and tedious and subject to a number of compensating errors (11, 16). HE STANDARD PROCEDURE
1230
ANALYTICAL CHEMISTRY
is formed; around pH 6 calcium hydroAlthough significant studies have gen phosphate, CaHP04.2H20 is prebeen made on the determination of cipitated; while above p H 6.9 hydrolcalcium in biological fluids by flame ysis occurs and the stable phase is a photometry (I@, this method has not won acceptance in the clinical laboramore or less pure material of colloidal dimensions having the crystal structure tory because of difficulties introduced of hydroxyapatite, Calo(P04)c(OH)2 by the presence of phosphate. Similarly, phosphate interferes: with t h r (IC). Although the solubility and the direct titration of calcium (I, S, 7 , IO, composition of hydroxyapatite are a function of the p H of the supernatant 23). However, because the (ethylenedinitri1o)tetraacetic acid (EDTA, Versolution and the solid-solution ratio, hydroxyapatite readily forms supersene) titration of calcium is more consaturated solutions (21). venient than the permanganate titration of oxalate, procedures have been Because of this and because calciuni and EDTA form a very stable complex developed whereby the phosphate is (fa),i t would seem possible to deterremoved on an anion exchange resin (12, 17) or by extraction into a 1mine calcium in phosphate solutions by merely changing the order of adding butanol-chloroform layer as phosphoreagents-Le., by adding exccss EDTA, molybdic acid (6). It has also been raising the pH, and then back titrating suggested that EDTA titration can with a standard calcium solution. replace the permanganate titration of The present work presents the results calcium oxalate (23). The difficulties observed in the EDTA of studies using this procedure for the deterniination of calcium in the presence titration of calcium phosphate solutions are due t o the specific and unusual of phosphate in materials containing magnesium and/or iron. Because of its properties of calcium phosphate itself. A t low pH’s moderately soluble calcium superiority oi7er murexide, calcFin was dihydrogen phosphate, C ~ ( H Z P O ~ ) ~ , H ~used O . as the metal indicator (6).
Other indicators such as Cal Ver 11, (Hach Chemical Co., iimes, Iowa) or Cal Red (Scientific Service Co., Dallas 21, Tex.) may also be used for the determination of calcium in strongly alkaline solutions. REAGENTS
Indicators. Calcein, prepared b y dissolving 2 grams of t h e indicator (G. Frederick Smith Chemical Co., Columbus, Ohio) in 25 ml. of 1M sodium hydroxide and diluting t o 100 ml. with distilled water. T h e solution was stored in a polyethylene bottle. Eriochrome Black T, prepared by dissolving 0.5 gram of the material (Fisher Scientific Co.) in 100 ml. of alcohol. The solution n a s stored in an amber bottle and kept in the refrigerator. Standard Calcium Solution. A solution containing 2.004 mg. of calcium per ml. (0.05OOM) l\-as prepared b y dissolving 5.005 grams of magnesiumfree, oven-dried calcium carbonate in 15 ml. of concentrated hydrochloric acid a n d diluting t o 1 liter n i t h distilled water. A O.OIOO1lf calcium solution \{as prepared b y making a fivefold dilution of the initial solution. The magnesium-free calcium carbonate was prepared by reprecipitation of Fisher's analytical grade calcium carbonate. EDTA Reagents. A solution having a titer of 1.1 mg. of calcium per nil. (0.029M) was prepared b y dissolving 17 grams of reagent grade disodium dihydrogen ethylenedianiinetetr aacetate (Fisher Scientific Co.) in 2 liters of distilled water. This solution was standardized (6) using calcein indicator. Identical results were obtained in the standardization of the E D T A reagent using the standard calcium solution and Iceland Spar. A second E D T A solution containing 0.133 gram of magnesium chloride, hIgCl? fiHzO, per liter of solution was prepared in the same ~ a and y standardized (8) using the Eriochrome Black T indicator. Ammonia-Ammonium Chloride Buffer. This n as prepared by mixing 67..5 grains of C.P. ainmoniuni chloiidv nitli 570 ml. of concentrated aqueous ammonia and diluting t o 1 liter. T h e buffer was stored in a polyethylene bottle. Sodium Hydroxide. ;1 1.5N solution nns prepared from C.P. sodium hydroxide pcllets and storcd in n polyethylene bottle. Phosphate, Magnesium, and Iron Solutions. il solution containing 1 mg. of phosphate per ml. n as prepared by dissolving 1.84 grams of anhydrous C.P. dipotassium hydrogen phosphate in 1 litcr of distilled a a t c r . KO further analysis was made. d solution containing 0.5 mg. of magnesium per nil. was prepared b y dissolving 1.4,5 gram3 of oren-dried C.P. magnesium cnrbonate in 10 ml. of concentrated hydrochloric acid a n d diluting t o 1 liter with distilled water.
This solution was standardized against the E D T A reagent containing magnesium chloride in a buffered ammoniaammonium chloride solution using Eriochrome Black T indicator. A 1 to 10 dilution of the magnesium solution was also prepared. An iron solution was prepared by dissolving ferric perchlorate in perchloric acid. After the iron content was determined in the usual way with a standard permanganate reagent, the solution was diluted to give a final solution containing 0.55 mg. of iron per ml. in 0.45M perchloric acid. GENERAL PROCEDURE
A 15 to 20% excess of standard E D T A solution was added to dilute hydrochloric acid solutions containing calcium and phosphate. The pH was raised to approximately 13 by the slow addition of 5 ml. of 4ilf sodium hydroxide. Then 2 drops of calcein indicator was added and the excess E D T A was back titrated slowly with the dilute standard calcium chloride solution until the solution became yellow-green. Finally, E D T A was again added until the pink-brown calcein end point was observed. End Point Illumination. Diehl and Ellingboe (6) recommend mixing charcoal with the calcein indicator or performing the titration in diffuse light, They also caution against titrating under a fluorescent lamp. Good end points mere obtained in this laboratory under lighting conditions whosc intensities at the norkbench nere less than those found in column tn-o of Table I. These variations are due to the differences in the rclative energies of the light sources at 500 nip (the rvcitation wave length for fluorescent solutions of calcein as determined by a Photovolt Corp. RZodt.1 54 fluoresecJnce meter and Corning interference filters) and the effect of light with different spectrnldistribution curves on color perception (9). However, the end point can he determined under normal lighting conditions either ivith the aid of a spot plate or b y performinq thc titration in a n-hite,
opaque polyethylene beaker one fourth to one third full. Most of the work reported here was performed with polyethylene heakers and magnetic stirring. RESULTS
Calcium in Presence of Phosphate a n d Magnesium. T h e results of calcium determinations in t h e presence of varying amounts of phosphate using t h e calcein indirator are given in Table 11. The direct titration it h e proccdure of Diehl and Ellingboe
Table I. Maximum Illumination for Observing Calcein End Point Illupinanre, Foot Candlesu Light Source Maximum N o r G l b Diffuse daylight 5 15-26 Incandescent 12 17-20 White fluorescent 17.5 10-45c a Measured at workbench with General Electric light meter. b Found in this lahoratory. because of design of fixtures and reflectors the illumination diie t o fluorescent lighting is greatlp dependent upon relative position of surface being measured. Table II. Calcium Determination in Presence of Phosphate (Each mmple contained 20.06 mg. oi citlciiim, as calcium chloride.) phosphate, Calcium Foiind, Me." hlg. Direct Indirect 13 1c) 87 20 06 21 19 91 20 06 32 19 81 20 08 40 19 90 48 19 88 ?O 10 20 10 64 1ct 80 70 19 91 80 19 94 20 06 96 20 06 5 Average of three determinations
Table 111.
Calcium Determination in Presence of Phosphate and Magnesium Calcium Found, _____ .- Present, Mg. Phosphate Magriesium bI$" Calcium 10 08< 10 03 11.5 0 0.025b 10.03 0.050 10.08' 0 125 10 08C 0 250 9 98d 0 50 10.03d "50 "0.0ii
:10.If
5.00 0 .025e
0.050 0.125
10 O l d 20,06.' 20.10" 20,080
0.50 1.25 2.50
5.00
20.04d 20. O l d 20.OY 20.12d
10.00
20.01d
Average of three determinations. * Total calcium and magnesium expressed as calcium equals 10.07 grams. c Clear solution. d 'Turbid solution. e Total calcium and magnesium expressed as calcium equals 20.10 grams. a
VOL. 31, NO. 7, JULY 1959
* 1231
( 6 ) , wherein the pH is first raised by tlic addition of sodium hydroxide and then indicator and EDTA reagent arc added. The direct titration method gave low results which were nearly independent of tlir phosphate concentration; precipitation occurred in each solution. The indirect titration method gave satisfactory results, with no 1)rwipitation and a sharp end point.
Table IV.
Calcium Content of Bones and Teeth Ca, yo .Ish weight ________ Accepted
Material Phlaiiges (33) Femurs (33) Humeri (33) Deciduous teeth (711
Foundu 39.13 39.09 39.04
values (130)
38 4
;3 i
,323 3
-4verage of four determinations on each bone or tooth. Table V. Calcium Content of Urine Sample Calcium, M g . Treatment Added Present Founda Ashed at 600" 3 ,323 Sulfuric acidpersulfate .3 44 0 68 4 12 4 18 1 21 4 65 4 62 1 99 5 43 5 33 2 61 6 05 5 96
.Iverage of four determinations.
Titration results for knoir n amounts of calcium in the presence of both phosphate and small amounts of magnesium are given in Table 111. In these experiments the calcium-phosphate weight ratio approximates that found in bone. When the amount of magnesium present is small, the calcium results are slightly high. K h e n sufficient magnesium is present to form a precipitate, the results tend to be slightly low, probably duc to the coprecipitation of calcium n i t h magnesium. -4s expected, the accuracy of the procedure increases with increasing calcium content. Similar low results for calcium in the presence of large amounts of magnesium were observed by Diehl and Ellingboe (6). The indirect E D T B titration was used for the determination of calcium in 99 human adult bones and 71 teeth from children. The methods of sampling and the biological significance of the data mill be presented elsewhere. The results expressed in per cent calcium of ashed material (Table IV) are in good agreement with data obtained by other observers. I n every case the ashed bone sample or ashed tooth was dissolved in hydrochloric acid and diluted to volume. dliquots of these solutions were taken for analysis. A 0.96 to 0.98 correlation was found between the calcium analysis of the individual bones by the indirect titration method and the mineral content of the same bones as evaluated by the tapered aluminum wedge technique @ , I @ .
Determination of Calcium in Presence of Phosphate and Iron Present, Mg. Calcium Found, Calcium Phosphate Iron Mg." 20.06 30 00 0 55 20 00 1 10 19 94 2.78 19 77 0 55 20 07b 1.11 20,OOb 2.75 19.92b Average of three determinations; all solutions turbid at end point. b Sodium hydroxide used to adjust pH contained 5 grams of sodium cyanide per 100 ml. of solution. Table Vi.
Table VII. Coprecipitation of Calcium with iron Hydroxide (Each solution contains 2.75 mg. of iron and 30 mg. of phosphate.) Excess calcium ~ ~ l Calcium ~ i , Found, ~ ~ Present, ina BackMg.b .\lp Titration, Centrif3Ig. ngate Residue 20 Ub
c
0 08 0 4 0 8
20 06 20 02 19 81 19 41
0 0 0 0
00 07 20 58
.\mount of calcium addctl piior to removal of iron hydroxide by centrifuging. Average of duplicate determinations. c Solution centrifuged just hefore back titration was completed.
1232
ANALYTICAL CHEMISTRY
The application of the indirect titration to the determination of calcium in urine is shown in Table V. The results were obtained on 10-ml. aliquots of a 24-hour sample of normal adult urine collected in a polyethylene bottle containing 10 ml. of concentrated hydrochloric acid. The ashed samples were dissolved in 5 ml. of dilute hydrochloric acid and analyzed in the usual way using a polyethylene beaker. The end point was detected with the aid of a spot plate. The other samples were decolorized by heating carefully for 20 to 30 minutes with 0.5 ml. of concentrated sulfuric acid and 1 gram of ammonium persulfate. Heating was discontinued when the samples were a straw yellow
and a precipitate formed. The solutions were cooled, a n excess of EDTA reagent was added, and the p H was adjusted by the addition of 10 ml. of 4M sodium hydroxide. The analysis was finished in the usual way. When the decolorized urine samples were made alkaline, the color deepened and the end point was difficult to detect. The analyses were performed using 0.01123M EDTA reagent and a 0.0100M calcium solution.
Calcium in Presence of Phosphate and Iron. The results of a series of experiments to determine the effect of iron upon t h e indirect titration foi calcium a r r given in Table VI. I n each eapeiiment the alkaline solutions containing excess E D T A \Tere clear and amber. Upon back titrating n i t h t h e standard calcium chloride solution, iron hydroxide precipitated. The back titration was continued until a permanent yellow-green cast could be observed and then E D T A was again added. The end point was difficult to detect due to the color of the iron hydroxide and to fading. Reproducible results were obtained by observing the formation of a permanent pink color along the walls of a polyethylene beaker. The addition of sodium cyanide decreased the error by a factor of 0.5, but it did not improve the end point. I n a n attempt to improve the end point, the iron hydroxide was removed by centrifuging. However, calcium coprecipitated with the iron. The solutions should be centrifuged just before the end point is reached (Table VII). These experiments were performed by mixing 10.01 ml. of 0.0500.M calcium chloride and 25 ml. of 0.02861M EDTA together with the iron and phosphate. Exactly 0.0100.1~calcium chloride was used in the back titration. After centrifugation, the calcium content of the residue was determined by dissolving it in dilute hydrochloric acid, precipitating the iron as the hydroxide by homogeneous precipitation using urea, recentrifuging, and analyzing the centrifugate by a direct titration iyith EDTA. Experiments were also carried out on a sample of XBS phosphate rock (56b). The rock was first dissolved in 5 to 10 ml. of concentrated hydrochloric acid on a steam bath. After several hours the solution was diluted and the calcium was determined by indirect titration. No attempt was made to bake out or otherwise separate the silica. I n some determinations the end point was determined by observing the formation of a pink color along the wall of a polyethylene beaker. I n other determinations the solutions were centrifuged just before back titration of the excess EDTA was complete, and the final end point was determined with the centrifugate in a polyethylene beaker.
~ l the ] determinat,ions were
lvith
diffuse daylight illumination. The rock contained 44.06% calcium oxide. ,4n average of four determinations each by the indirect and cent& fuging methods gave 44.17 0 . 0 ~and 44.10 z!= 0.020/,, respectively.
*
LITERATURE CITED
(1) Bete, J . D., Koll, C. .4.,J . A m . Water Works Assoc. 42, 49 (1950). (2) Broxn, W. N., Jr., Birtely, W. B., Rev. Sci. Instr. 22, 67 (1951). (3) Cheng, K. L., Bray, R. H., Soil Sei. 72,449 (1951). (4) Clark, E. P., Collip, J. B., J . Biol. Chem. 6 3 , 461 (1925). (5) Collier, R. E., Chemist dnalyst 43, 41 (1954).
w.,
(6) Diehl, H., Ehigboe, J., ANAL.CHEX. 28,882 (1956). ( 7 ) Diehl, H., Goetz, C. A., Hach, C. O., J . A ~ Water . works A ~ 42,~ 40 ~ (1950). (8) Diehl, H., Smith, G. F., “Quantitative Analysis,” p. 343, TViley, Xew 1-ork, 1952. (9) Evans, R. >I., “Introduction to Color,” Chap. 111, Wiley, Sen- York, 1948. (10) Fales, F. W., J . B i d . Chem. 204, 577 (1953). E. F., (11) .Hillebrand, W.F., Luntlell, Bright, H. A,, Hoffman, J . I., .Applied Inorganic Analysis,” 2nd ed.! p. 621, New York, 1953. (12) Jenness, R., AXAL. CHEN. 2 5 , 966 (1953). (13) Kramer, B., Tisdal F. F., J . Biol. Chem. 47, 475 (1921j . (14) Levinskas, G. J., Neuman, K. F., J . Phys. Chem. 59, 164 (19551.
(15) AlcFarland, Science 119, 810 (1954). (16) ZIIcIntyre, I., Biochem. J . 67, 164 ~ (1957). . (17) Mason, A. C., Analyst 77, 520 ( 1952). (18) Schwarxenbach, G., Ackermann, H., Helv. Chim. Acta 30, 1798 (1947). (19) Sobel, A. E., Skersky, L., J . Biol. Chem. 122, 665 (1938). (20) Spector, W. S., ed., “Handbook of Biological Data,” W. B. Saunders, Philadelphia, 1956. (21) Strates, B. S., Neuman, W. F., Levinskas, G. I,., J . Phys. Chem.61,279 (1057). (22) Wang, C. C., J . Riol. Chem. 1 1 1 , 143 (1035). (23) Wilson: A. E., ASAL. CIIEM. 22, 1571 (1950).
9.
RECEIVELI for review October 13. 19\58. A 4 r c ~ p t elrarch d 20, 1959.
Kinetics of the Oxidation of Bromide Ion in Fused Alkali Nitrate Solutions of Dichromate F. R. DUKE and M.
L.
IVERSON
Institute for Atomic Research and Department o f Chemistry, lowa Sfate College, Ames, lowa ,The oxidation of bromide ion b y dichromate in fused sodium nitratepotassium nitrate eutectic proceeds through the steps: Cr20;z NO, $ NO; 2 CrOa2 and BrNO’,
+
+
+
NO: f Br.
Lead ion is added to precipitate chromate ion and drive the first equilibrium equation. Corrections must b e made for the complexation of lead ion b y dichromate and bromide ions.
is not ordinarily an active oxidant in aqueous solution; rather, the acid chromate ion (HCrO;) is gcnerally responsible for oxidations ( 6 ) . I n fuscd salt media, the equilibrium is not possible. Thus, i t is of interest to find whether or not dichromate ion will react directly with a reductant in these media. I n fuscd alkali nitrates, there is a. rapid equilibrium of the sort ( 1 ): ICHROMIATE ION
Cr?O;’
+ NO;
e NO:
maintained = 1’ C. by a Brown controller.
nere prepared in fused sodiuiii iiitratepotassium nitrate eutectic. To these were added solutions of lead nitrate and potassium bromide in thc same solvent, the lead being a precipitant for the chromate, thereby controlling the concentration of the latter. The concentrations of the reactants were varied from one run to another. The reaction yessel mas made of borosilicate glass in the shape of a large test tube fitted with a ground glass stopper carrying an inlet tube, which reached t o the bottom, and an outlet tubp. Preheated purified nitrogen was passed through the samplr a t the rate of 0.21 liters per minute and the evolved bromine was absorbed in an acidified aqueous solution containing sulfite and silver ions. The escess silver n as determined by a Tolhard titration as n function of time. A11 solutions, the gas preheater, and the reaction vessel were immersed in a constant temperature bath of fused alkali nitrates, with the tcniperaturc,
RESULTS AND DISCUSSION
I n the interpretation of the result>, the data were differentiated by making a plot of bromide ion us. time and dran ing normals to the smoothed curve a t various points; the values of -d[Br -11 dt were taken from the slopes of perpendiculars to the normals a t varioui reactant concentrations. The values ot -d[Br-] clt and the corresponding total arid uncomplexed concentrations of rcactants are shoivn in Table I. If nitroniuni ion is assumed to he the oxidant in the solutions, then
The probability of the reaction being first order in nitronium ion is large,
+ 2 Cr0;’ (1)
The nitroniuni ion might be expected to oxidize bromide ion. It would be expected that variations in chromate ion would affect the rate of the reaction of nitroniuni with bromide, but that the direct reaction of dichromate with broniidc would not be influenced b y chromate. EXPERIMENTAL
Table I. Rate Data for Dichromate-Bromide Reaction in Fused Nitrates
(kK/K2,,,calrulated according to I?quation 7. n = 1)
0 00G6
0 0027
0 0053 0 0063 0 0018 0 0021 0 0042
0 067
0 059 0 064 0 031 0 022 0 023 0 027
0 00‘36 0 0057 0 0082 0 020.1
0 0160 0 0163 0 0181
0 0025 0 0010 0 0019
0 0039 0 0010 0 0014 0 0018
0 056 0 055
0 0 0 0 0
055 020 015 016 018
0 0026 0 0017
0 0026 0 010
0 0095
0 0093
0 010
0 00033
0 00023
0 00085 0 00074 0 00014 0 00019 0 00048
6 6 6 5
X X X X
10‘ 10‘ lo1
10‘ 6 X 104 6 X lo4 6 X 10‘
Solutions of potassium dichromate VOL. 31, NO.
7,JULY 1959
1233
