Discussion. T h e total shaking time beemh to depend not only on the particle size of the sample, but also on the type of metal under study. It may depend on the chemical affinity of the metal toward oxygen. Pure copper metal granules, 0.05 ;gram of approximately 0.5-mm. mech size, required about 2.5 hours of .haking, as compared to 50 minutes for tin. The longer time needed for complete cold combustion of
the copper may be due t’o its nobility.
LITERATURE CITED
Aisample consisting of 0.01 gram of gold
(1) Habashy, M. G., ANAL.CHEM.33, 586 (1961). (2) Ibid., 34, 1015 (1962). (3) Hopkins, D. W., “Physical Chemistry and Metal Extraction,” J. Garnet Miller, London, 1954. ( 4 ) hfellor, J. W., Parks, G. I)., “Modern Inorganir Chemistry,” Longmans, Green, London, 1956.
showed no signs of oxidation, even after 4 hours of shaking. A study is being continued to define the actual relation between the time needed for t,he quantit’ative cold combustion of a certain weight of metal and its chemical affinit,y towards oxygen. Such a relation may allow the direct determination of t’he chemical affinities of amalgable metals by the cold combust,imeter.
RECEIVED for review December 17. 1962. Resubmitted April 2, 1064. Accepted April 2, 1964.
Determination of Traces of Alkaline Earths in Alkali Halides by Spectrophotometric Titration in the Ultraviolet E. P. PARRY
and G. W. DOLLMAN’
North American Aviation Science Center, A Division o f North American Aviation, Inc., Canoga Park, Calif.*
b Titration of traces of alkaline earth in the presence of a very large excess of alkali halide using the large absorbance in the ultrciviolet (222 mp) of the alkali EDTA complex is discussed. As little as 0.5 p.p.m. calcium can b e determinted in any of the alkali halides using ai 0.5-gram sample. The titration will determine any of the alkaline earth ions or their sum if a mixture is present. Titrations of barium, magnesium, and strontium in lithium chloride and Ibarium in sodium chloride cannot b e done accurately in the lower concentration ranges because of the preferential ,Formation of the alkali EDTA complex before the end point. The quantitatiive nature of this effect is discussed.
I
STUDIES of the mechanical and ipectroscopic proprties of alkaline earth doped alkali halides, the accurate determination of the amount of alkaline earth ion present in the crhstal 15 of crmiiderable importance. I3ec auie of the need either to dtlteimine very loa le1eli of alkaline earth or to analhze \ e q -mall \ample\, the uwal \i.ual titration procedures 3 re not adequate Other l)rocedure< such as the wren proposed neutron artil ation anal> 1)rocediire (8) or iome qwtrophotometric techniques ( 4 , 6) for microgram amounts of calcium either are not applicable or have a number of diqadvantage%when applied to the problpm a t hand. -1i ~ ~ e c t r o ~ ~ l i o t o m emicrotric titration of calrinm using an indicator has heen deitribed (I), but the pro-
s
’ Present address, Space and Guidance Control Division, Lytton Industries, 5500 Canoga ilvenue, Woodland Hills, Calif, Now at Thousand Oaks. Calif
cedure would only determine a minimum
of 40 p.p.m. calcium in potassium chloride. Yo study wa’ made of the effect of large amounts of alkali metal ions on the titration. For the determination of other alkaline eart,hs, more serious problems can arise, although a polarographic met’hod has recently been described (10) for determination of barium in potassium iodide. Plumb, Martell and 13ersn-orth ( 7 ) first showed that (ethy1enedinitrilo)tetraacetic acid (EDT.l) absorbs more strongly in the ultraviolet than some of its metal complexes. Sweetser and Bricker (9) applied this ohservat,ion to the spectrophotometric titration of several metals. The present work use5 the same ultraviolet8m d point. but (>onsiders the determination of traces of various alkaline earth ions in the presence of a large excess of various alkali halides. 21s little as 0.5 p.p.m. calcium can be det’ermined in any of the alkali halides in a titration requiring about 20 minutes. However, the applic~atioriof the procedure to some of the other alkaline earth:: in certain alkali halides is 3everely limited. While no effort was made to eliminate interference of transition metal ions (since they were of no significance in these experiment,s),chelat,eion exchange resins could undoubt,edly he employed for their separation if necessary. EXPERIMENTAL
Reagents and Apparatus. Reagent grade chemicalq were uied without further purification. .4 1 m J i EDT.l solution was prepared by disqolving 0 375 grams of dihydrogen diiodium ethylenediamine tetraacetate in nater and diluting to 1
liter. The solution was standardized against a standard calcium solution. 1.1.f ammonium chloride-ammonia buffer of pH 10 was prepared by diluting 5 . 4 grams of SH,C1 and 75 ml. of ammonia (sp. gr. 0.90) to 100 ml. h standard (10-3.1f) calcium solution was prepared by dissolving an accurately weighed amount (0.1 gram) of calcium carbonate solution in a minimum amount of 3.11 HC1 and diluting to 1 liter. A Car. Model 14 spectrophotometer was used for absorbance measurements. One-centimeter quartz cells of about 3.5-ml. volume were used as the titration cells, and the t’itrant was added with a 2-ml. capacity micrometer buret (Cole Palmer Instrument and Equipment Co., Cat. S o . 7486). The contents of the cell were stirred by a fine stream of gas bubbles after each titrant addition. Procedure. The solubilities of sodium chloride, potassium chloride, arid lithium chloride a t 0” C. are 3 5 . 7 grams:100 nil., 2 7 . 6 grams/100 ml., and 67 grams; 100 ml., respectively. I n order to maximize the sensitivity, solution.: near saturation should br t’aken. However, some mlutions that are too conccintrated give rise to later. For difficulties, a. diwu..ed this reason, a sample size of 0 . 5 gram in 2.0-nil. solution (25 grams/ 100 rnl.) wa.: t,akrn as opt’imum for all alkali halides. If sufficient sample is available, accurately weigh approximately 1 , 2 5 grams of sample into a 5-rnl. volumetric flask, add 10 drops of buffer solution, dilute to volume) and shake well. Using a volumetric pipet transfer 2 . 0 ml. of solution to a I-cm. quartz cuvet (volume of cuvet -3.5 ml.) and Illace in the cell compartment of the spectrophotometer. Read the absorbance a t 222 mp. Add 1 m J i EDT,I in small increment’s reading the absorbance after each reagent addition. The change in absorbance should be small until after VOL. 3 6 , NO. 9, AUGUST 1964
1783
the end point, and then a sharp rise will occur with increasing amounts of reagent. h typical titration is shown in Figure 1. The data should be corrected for dilution before plotting. .I blank should be prepared in an identical manner but without the sample, and the titration volume obtained should be subtracted from the sample end point volume. The sample size should be adjusted so that the titration does not require more than 1 . 3 ml. of reagent, because it becomes difficult to stir the cell contents when it contains too much solution. The use of 0 . 5 gram sample in 2 . 0 ml. of solution permits the facile determination of 0 , 5 to 25 1i.p.m. calcium in various materials by this procedure. If the sample contains as much as 100 p.p.m. calcium, an accurate titration can be made on as little as 50 mg. of sample using the normal technique. RESULTS A N D DISCUSSION
The absorption of EDTA in the ultraviolet and the shift of the absorption band to lower wavelength by complex formation was pointed out by Plumb, Martell, and Bersworth ( 7 ) . We confirm their observation of not being able to reach the maximum of the absorption band since it is too far in the ultraviolet. The sensitivity of the titration inFor creases with increasing pH. example, the absorbance of 1 m M EDTA solution without any alkaline earth salt present is 1.56 a t a pH of 10.4, 1.40 a t a p H of 10.0, and 1.25 at a pH of 9.40. However, if magnesium is to be titrated, the pH should be less than 10.5 in order that Mg(OH)2 not be precipitated. The optimum pH for the titration is therefore taken as between a pH of 10.0 and 10.5. I t is obvious that sufficient buffer should be used to keep the pH constant during the titration.
Table I. Titration of Alkaline Earth Ions in Various Alkali Halides
Alkali halide KC1 KC1 KC1 KC1 KC1 KC1 KC1 SaCla L1Cla
KC10 KC1. NaCla 1,iCIa -.-.
Alkaline Amount A.E. earth (.4.E.) added, ion used p,p.m. 0 Ca 1 0 Ca 2 5 Ca 5 0 Ca 10 0 Ca 15 0 Ca M g Me
ca Ba Ra
Mg
Ca - ..
+ Ca
+ Ca
5 0 10 0 9 7 30 0 60 0 14 7
.io
0
ANALYTICAL CHEMISTRY
0
!
I
I
1 I
01
02
03
0.4
VOLUME ( m l s ) EDTA;
Figure 1 . Spectrophotometric titration of in a sample of potassium chloride
Relat’ion of the band shift to stabilit’y of the complex is shown in Figure 2 , where a plot of xo.3 tis. log stability constant of the complex is given (h0.3 is the wavelength a t which the absorbance is 0.3 unit). It is clear that a definite separation exists between the wavelength where t,he alkali complex ions absorb and that where the alkaline earth complex ions absorb. There is also relatively little difference between the shift by a complex of relatively low stability (I3a+2) and that for a complex of high stability (Zn+*). The above observation suggests t,hat it would be difficult t,o determine individual alkaline earth ions in a mixt’ure but relatively easy to determine their sum. Table I shows some typical data. Figure 2 also suggests t’hat it should be possible to titrate barium ion in the presence of lithium chloride. However, results for this titration are always very low. This titratmionis subject to serious interference because the lithium EDT.l complex forms before the end point is reached. An equation can be derived which indicates the magnitude of this interference. By writing the equilibrium constants for the alkali and alkaline earth complexes and the electroneutralit’y and mass conservation equations, the following equation can be obtained.
9.3 p.p.m. calcium
13-
12-
1-
0
Total concentration of alkali ion = Formation constant for alkKfA ali-EDT-1 complex [ H f ] = Hydrogen ion concentration a t which titration is done p, p‘, and 8” = Aictivity coefficient ratios
M,
=
Total A.E. found, p.p.m. 0 28 1 4 2 8
where
10 2
b
5 6
15 5 10 9 30
4 2 2 8 1 61 0 14 4 50 5 27.4 8.8
Ca 25.0 LiCla LiClS Ca 10.0 a Corrected for titratable8 in alkali halide.
1784
I
Y
M a
K,E
K4
Concentration of alkaline earth complex = Total concentration of EDT.1 added = Total concentration of alkaline earth ion = ~[H+l p”KfEK4 = Formation conqtant for alkaline eartli-EDT.l complex = The fourth acid dissociation constant of EDT.l =
I n the derivation of this equation, we have restricted ourselves to a p H range where the formation of H 2 P 2 can be neglected (pH > 9) and have made one very reasonable assumpt,ion t’hat the concentration of alkali ion in the romplex is very small compared to the total concentration. The p terms represent the activity coefficient ratios which convert concentration constants from the ionic strength at which they were determined to the ionic strength of t h r niedium wed. Since 13 is difficult t o evaluate, Equat,iori 1 cannot be used to
PR
Figure 3. Plot of pR( ---log R) (See text) vs. log concentration of alkaline earth ion
predict accurately end point error, but it can indicate where an error might be expected. The term [H+]lp"K/EK, reflects the formation of the monohydrogen EDT.1 ion with decreasing pH. This will increase the extent of dissociation of the alkaline earth complex wit'h decreasing pH. Howewr, a t p H 2 10, this term is negligibly small [with reY)]for the lowest conspect to (M centrat.ions of alkaline earth titrated, and can be neglected in the calculat'ions. The term l:'p'KjE is also negligibly small and can be neglected under all condit,ions of the titration. The plus sign before the square root in Equation 1 should be used after t'he end point, and the negative sign should be used before the end point. In the absence of alkali metal ion and at p H = 10, this will reduce to b = Y before t,he end point, and to b = -11after t8heend point. If the equation ic; derived in t,he absence of alkali ion, i t is identical except it does not contain the term
+
MI
(+:)
This term represents the
interference of the alkali ion on the titration of the alkalinje earth. I n order to evaluate> the possibility of interference in various mixtures of alkaline earth ions in alkali halides, the
calculated at the equivalence point for 10 p.p.m. alkaline earth ion in the halide salt (assuming B to be 1). The per cent of the EDT?i which is in the alkali ion complex is also given. The results are given in Table 11. This indicates qualitatively the magnitude of the error which might be expected. The ratio is negligibly small for some mixtures, such as calcium in lithium chloride and magneiium in sodium chloride, and the data of Table I show that such titrations can be done. On the other hand, for bai*ium, magnesium (and strontium because its stability constant is nearly the same as magnesium) in lithium chloride, and for barium in sodium chloride, substantial errors will result in the titrations. Typical data are shown in Table 111. From Equation 1, it is apparent that
the error will decrease as the concentration of alkaline earth ion increases, as the concentration of alkali metal ion decreases, and as the ratio of formation constants of alkali metal to alkaline earth decreases. I n Figure 3, a plot is given of pR(-log R) us. the log concentration of alkaline earth ion. This plot indicates, for example, that 3160 p.p.m. barium ion in lithium halide should show approximately the same titration error as 3.16 p.p.m. magnesium in sodium halide. However, no quantitative relationships can be given since the end points are not well defined and the activity coefficient ratio ( p ) is not known. From inspection of Tables I and 11, it appears that if -log Risabout 2.25 to 2.50 and greater, the titration can be done reasonably accurately -.
p
26-
I
24-
iI
22-
20-
0
I
EL 0
1
2 1 MOLARITY SALT
4
5
6
Figure 5. Effect of alkali chloride concentration on absorbance of 1 mM EDTA solution; 1 -cm. cells
points from the calculated curve is caubed by an increase in absorptivity of the LiEDTX complex with dilution. T o demonstrate this change in absorptivity, Figure 5 shows a plot of absorbance of a 1 m M E D T A solution of pH 9.4 as a function of molarity of KC1, LiC1, and NaCl. An insignificant change is noted for potaisium chloride while sodium chloride shows a small change. However, a large change is
Table II. Interference of Alkali Ion on Titration of Alkaline Earth Ions
(for 10 p.p.m. A.E. in alkali halide at equiv. pt.)
90
I
O'
01
02
03
04
05
MLS EDTA (001
06
d7
08
E1
Figure 4. Comparison of calculated and experimental curves for the titration of 1 5 5 4 p.p.m. barium in lithium chloride; experimental points; solid curve is calculated
Using approximately 1500 p.p.m. barium in lithium chloride and titrating with 0.OlM EDT.4, a titration curve can be obtained and compared with that calculated from Equation 1. The absorptivities for the B a y p 2 complex and the LiY-3 complex were obtained from the beginning slope of the titration curve for the individual ions. In order to obtain p , the activity coefficient ratio, a value was determined by curve fitting to the experimental point at 0 . 2 ml. added. Using these parameters, the curve was calculated for the complete titration range and is shown in Figure 4. The fit with the experimental points a t the high end is only fair, but the interference of lithium EDT.1 complex before the end point is well demonstrated (compare with Figure 1). The deviation of the experimental
Mixture Ra Ca in LiCl 6.96 X Mgin NaCl 1 . 9 X 10-3 M a in LiCl 3.6 X B a i n NaCl 9.35 x 10-2 Ba in LiCl 1.73 a R = M I (K,A/KrE)/(M text).
+
EDTA as alkali EDTA complex Si1 6 23 35 81
Y ) (See
Table 111. Typical Titration of Interfering Mixtures
Alkali halide present LiCl LiCl LiCl NaCl LiCl LiCl LiCl LiCl LiCl
Alkaline earth ion titrated Mg
Mg w 3 Ba Ba Ba Ba Ba Ba
Amount A.E. added 10 30 100 44 200 400
1500 1,550 1650
Total A.E. ion found 5 8 18 4 92 5 35 0 None 112a 105Oa 12TOa
1100~
These end points are not reprodurible since no straight line portion is obtained for extra dation, and the end point is greatly &pendent on the number and position of points after the end point.
VOL. 36, NO. 9, AUGUST 1964
1785
observed for lithium chloride in a direction such that an increase in absorptivity would be observed with dilution of the alkali halide sample. The large increase in absorbance in going from no salt to 0.35df salt repto resents the change from free EDTA% LiEIITA complex. The absorbance change in going from 0.35Jf LiCl to 5.663f LiCl cannot be caused by a shift in equilibrium in the free EDT.1LiEDT.4 ratio with ionic atrengt'h, since there is essentially no free EDTA in any of t'he solutions. This is a t once apparent by considering the LiEDT.4 equilibrium expression. Thus a change in the equilibrium constant wit,h increasing ionic st'rength cannot be responsible for this behavior. It. is possible that this increase in absorbance with increasing salt content could be caused by an increase in refractive index of the solution. From dispersion theory, it can be shown that' t,he following expression is valid if no chemical changes occur in the absorbing species and if the damping factors and oscillator strengths of absorbing species are const'ant (3).
x ahsorpt'ivity = constant where n is t,he solution refractive index a t t'he wavelength in question.
n
Thus a decrease in refractive index with decreasing salt content would
necessitate an increase in absorptivity. Using a refractive index of water (5) of 1.39 and an extrapolated value for lithium chloride of 2.3 at 222 mp and values given for the specific gravity of lithium chloride solution (a), it is possible to obtain t,he values of the molar refraction, and from that, the index of refraction of various lithium chloride solutions a t 222 mp. Such calculations show t,hat the change in refractive index is much too small to account. for the large change in absorptivity. -4possible explanation for this change in absorptivity can be attributed to chemical changes in the solution. Thus additional complexes-i.e., Li2Y-2, LisY-, etc.,-could be formed with increasing lithium concentration, with such species having significantly different absoqhvities. While no proof of this is given, it appears consist'ent with t'he present data. N o drastic effect on the results for the titration in LiCl is caused by this increase in absorptivit,y upon dilution as is shown by Table I. However, the result's are riot' as reproducible and accurate as those obt,ained in the titrations done in KC1. This is undoubt'edly caused by a dependence of the slope of the line aft'er the end point upon the exact volumes at which the absorbance values are read. Results of Table I appear to be accurate within
10%. Better results could certainly be obtained by titrating with an EDTA solution in very pure 25% LiCl, so that no dilution is obtained in the titration. Such a technique should considerably improve the preciyion obtained. The data of Figure 5 would a h 0 suggest the possibility of titrating moderate amounts of lithium in potassium chloride using this technique. LITERATURE CITED
( 1 ) Chalmers, R. A., Analyst 79, 519 (l!54). ( 2 ) Handbook of Chemistry and Physics," 44th ed., p. 2118, Chemical
Rubber Publishing Co., Cleveland, Ohio, 1962. ( 3 ) Hansen, W. X., NAA Science Center, private communication. ( 4 ) Herrera-Lancina, M.,West, T. J., ANAL.CHEM.35, 2131 (1963). ( 5 ) Landolt-Bornstein, "Numerical Values and Functions," 6th ed., Vol. 2, Part 8, p. 566, Optical Constants, Springer Publishing Co., Berlin, 1962. ( 6 ) Satelson, Samuel, Penniall, Ralph, AKAL.CHEM.27, 434 (1955). ( 7 ) Plumb, R . C., Martell, A. E., Bersworth, F. C., J . Phys. Chem. 54, 1208 (1950). ( 8 ) Strelow, F. W. E., Stork, H., ANAL. CHEM.35. 1154 (1963). ( 9 ) Sweetsei, P. k'., Bricker, C. E., Ibid., 26, 195 (1954). (10) Whitnack, G. C., Hills, &I. E., J . Electroanal. Chern. 6 , 68 (1963).
RECEIVEDfor review March 3, 1964. Accepted June 1, 1964.
Formation of Lead Chloride Fluoride at Lower pH Values for Gravimetric Purposes RAYMOND A. BOURNIQUE and LIONEL H. DAHMER' Chemistry Department, Marquetfe University, 12 I7 W. Wisconsin Ave., Milwaukee, Wis. 53233 Formation of PbClF by modification of the Hoffman and Lundell procedure to permit lower pH values was accomplished with sodium formate or sodium chloroacetate in place of sodium acetate as buffering agent. The auxiliary precipitant, chloride ion, was added in an appreciably decreased amount of HCI only rather than as HCI plus NaCI. These conditions permitted good gravimetric and volumetric recoveries of fluoride and satisfactory values of the Pb: CI ratio when the pH was maintained at 1.8 to 2.0. Precipitations were carried out in polythene beakers to avoid attack of glass by HF. At pH values below 1.8, results were definitely low while above pH 2.0 appreciably high values were obtained. Sodium carbonate in amounts as large as 2.0 grams did not interfere at pH
1786
ANALYTICAL CHEMISTRY
2.0. Hence the procedure seems suitable for gravimetric determination of fluorine in organic samples after sodium peroxide fusion. The presence of carbonate in such amounts causes slightly low results.
s
1911 when Starck (10) reported the use of lead chloride fluoride for the gravimetric determination of fluorine, a number of volumetric modifications (3, 6 , 6, 8, 12) of the procedure have been proposed. Most of these involve the titration of the chloride in the dissolved precipitate by a Volhard technique. At the high acidity involved in this titration, anions of weak acids such as carbonate generally do not interfere. In gravimetric work, however, lead salts of wch anions will often be found in the lead chloride fluoride. INCE
-1recent volumetric approach is that of Vrestal and coworkers (12) in which an excess of lead chloride is used to form the precipitate in a neutral solution, follon ed by complexometric titration of the excesb reagent. The procedures baied on titration of chloride in the precipitate have she\\ n appreciable deviations in results obtained. I n 1949, Kaufman (7') reported that insufficient control of pH may be responsible for this. I n studies of the Hoffman and Lundell (4,s)method, he concluded that the pH should be maintained between 4.6 and 4.7 for satisfactory results. Variations obtained with different volumetric PbClF methods were obPresent address: Department of Chemistry, Iowa State Univereity, Ames, Ion-a.
