1410
ANALYTICAL CHEMISTRY
calculated and known percent:tges of ortho-; the remaining columns give the corresponding data for the meta isomer. Where several observed values are given for a particular mixture, they represent the results of repeated determinations made at other times over a period of several weeks. ‘DISCUSSION
The largest source of error in the determination is the uncertainty in absorbancy due to the noise-to-signal ratio. The noise lev61 was variable but generally amounted, to about *l mm. (0.001 microvolt) or less 0 1 1 t,he recorder chart, corresponding t o approximately 0.5% of the total deflection. For o-cresol a t 5% roncentration, this corresponds to *0.40/;,error; for m-cresol a t t.he same Concentration, the anticipated error would be +O.S%. The average deviation of all the results is *0.3y0 ortho and *0.57, meta. These values are somewhat more favorable than the calculated errors, probably because two or more successive wave-length scans were always averaged for each determination. Samples of purified p-cresol prepared in the different ways previously outlined were compared on the basis of their spectra in the region 12 t o 1 4 ~ ; these were found to he indistinguishable a t concentrations of about 1 M.
The results of this study show that an accuracy of approximately *0.3% for o- and t 0 . 5 7 , for m-cresol is possible in the low conrentration region with the infrared spectrophotometric method described. With the use of absorption cells of greater thickness and a split-beam type spectrophotometer, a considerable improvement in precision should be attainable. Unfortunately, these were not available. The present method does not appear suitable as an analytical technique when the purity of the p-cresol is above 99%. Preliminary experiments have indicated that the freezing point depression offers a more sensitive analytical method for t o t d impurities in the region from 0 to 2% impurity. LITERATURE CITED
(1) Friedei, R. A., Pierce, L., and McGovern, J. J., A 4 ~ a C ~~ .M .22,. 418 (1950). ( 2 ) Sutherlapd, G. B., and Willis, N. A . , Trans. Faraday SOC., 41, 181 (1945). (3) Whiffen, D. H., and Thompson, H. N‘,, J, Chem. Soc., 1945, 268. (4) Wootfolk, E. O., Orchin, M., and Dull, M. F., IND. ENG.CHEM., 42, 552 (1950).
RECEIVEDMay 1 , 1950.
High Frequency Titrations Mercurimetric Determination of Chloride W. J. BLAEDEL AND €1. V. MALMSTADT University of Wisconsin, Madison, Wis. Chloride may be determined in acidic solution by direct titration with 0.01 kf mercuric nitrate using the high frequency titrimeter with a precision corresponding to ahout 0.03 ml. of mercuric nitrate. The end point agrees well with the equivalence point over a considerable range of conditions, allowing a simpler procedure than is possible when chemical indicators are used to establish the end point. For this titration, the high frequency titrimeter appears superior to the potentiometric or conductometric methods for establishing the end point.
C
HLORIDE may be determined volumetl’ically by titration
with standard mercuric nitrate according to the reaction : H g + + 2C1- +HgCIS (1) .Is stated by Kolthoff and Sandell ( J ) , this determination is of considerable practical importance because it allows direct determination of chloride in acid medium even at great dilutions. Jn:ismuch as there do not exist many such metho.ds, it seems worth while to attempt improvement of the mercurimetric procedure. Any of several substances may be used a? indicators for the rnercurimetric titration of chloride. Perhaps the most careful investigation of the procedure is that by Roberts (6), in which 1,5tliphenylcarbohydrazide is used a6 an indicator. The difficulty with thiE indicator (and with others) is that the end point is considerably different from the equivalence point and that large blanks are therefore required. The blank-which is as high a8 0.5 ml. of 0.01 M mercuric nitrate in some instances ( 1 , 6)-is highly dependent on the conditions of the titration, such as: mercuric chloride concentration at the end point; indicator concentration ; acidity; and ionic strength. Close adherence to carefully @electedconditions is necessary for accuracy, and this makes for .some inflexibility and inconvenience in titration procedure. Particularly troublesome is the dependence of the blank on the amount of sought-for substance-i.e., chloride. Kolthoff and Sandell (3) recommend use of sodium nitroprus.
+
side as an indicator, but unpublished work (1) has shown this inferior to 1,5-diphenylcarbohydrazide. Among other things, nitroprusside is unstable, decomposing to cyanide, which reacts with mercuric ion. Potassium iodate and potassium periodate are superior to sodium nitroprusside, but inferior to 1,Sdiphenylcarbohydrazide (1). These difficulties are great enough to prevent widespread use of the mercurimetric procedure. Many of the difficulties are due primarily to inadequacy of the available chemical indicators. By u&2of an instrumental method to establish the end point, these difficulties may be circumvented. I n this paper, a comparison is made among the potentiometric, conductometric, and high frequency methods of establishing the end point for the mercurimetric determination of chloride. The limiting conditions, together with the relative advantages, of the high frequency procedure are given. COMPARISON O F POTENTIOMETRIC, CONDUCTOMETRIC, AND HIGH FREQUENCY TITRATION PROCEDURES
No extensive study has been madeof the possibility of establishing the end point potentiometrically for the mercurimetric titration of chloride. The prospect6 for doing so do not seem good. Silver-silver chloride or calomel electrodes are not stable in solutions containing mercuric ion ( 1 , 4 ) . Muller and Aarflot ( 5 )claim
1411
V O L U M E 22, N O . 11, N O V E M B E R 1 9 5 0
in the potentiometric case. I n obtaining the data, an Industrial Instruments bridge (Model RC-1B) was used, and the titratiori vessel was thermostated a t 23" * 0.1 O C. Inspection of the conductometric titration curve (curve H, Figure 1 ) shows that the end point is not sharp. Three factors contribute to this lack of sharpness: aa w o n WW
wo
20
aa
It
vv W
0;;f W
a,. &a 3
!a
UJW
m l 6
7
8 9 10 II 12 13 14 ML. MERCURIC PERCHLORATE.
13
Figure 1. Comparison of Potentiometric, Conductometric, and High Frequency Titration Procedures in Mercurimetric Determination of Chloride Conditions: 20 ml. of'O.01 M sodium chloride titrated with M mercuric perchlorate in end point volume of 80 ml. End points: Theoretical, 10.86 ml.; potentiometric, 10.65 ml. at 375 mv.; Conductometric, 10.70 ml. at 0.00325 reciprocal ohms; high frequency, 10.84 ml. Scale: One ordinate scale unit equals 100 cyelea/sec., or 100 mv., or 0.0003 reciprocal ohms
0.00920
that chloride may be titrated potentiometrically with mercuric perchlorate, using mercury and normal calomel electrodes, I)utdetails of the procedure are not given. I t was not possible to reprdduce the work of Muller and Aarflot in this laboratory. B y adding metallic mercury t o solutions of mercuric perchlorate containing chloride, mercurous chloride was always formed. This was true whether chloride was present in low concentration or in excess of the mercuric perchlorate and also whether the solution was only slightly acid (pH 3 to 4 ) , or highly acid (pH 1). The same results were obtained with mercuric nitrate instead of mercuric perchlorate. This work, in agreement wit,h the standard potentials involved, showed the mercury electrode t o be unstable in solutions containing mercuric ion because of thereaction: HgCI? Hg +IIg,Cl, (2)
+
1. The presence oi perchloric acid in the standard merc*uric nitrate diminishes the difference i n slopes before and alter thr chquivalence point. 2. Curvature in the wings of thc titration curve niaki,- tor difficulty in locating the equivalence point. .4lthough thi, c > i i r l poirit of curve R, Figure 1: agrees imsonably wit,h the theoi,t.!ic,:il, the curve could have been just, as validly drawn to give itti t x n t l point diRering from this by 0.1 to 0.2 ml. I n cases such as t l i i < , the spacing of points and personal judgment influence location (I!' the end point greatly. 3. Curvature in the region of the equivalence point, c l i i c ~ t o formation of mercurous chloride ions, Ppreads out the e r d lloitit. This cannot be serious, .however, since the high frequewy c w l point is subject to the same error but is quite sharp. The high frequeiicy curve-curve C, Figure 1-for the s : i n i ( A titration as described for the potent,iometrie case, possesses :I sharper end point than the potent,iomet;ric or conductomrt1,ic curves and appearp su elior as far as the mercurimetric titration of chloride is concernel 'The factors which cause spreading of thc c'onductometric end point are minimized by use of the differrnti:Il titration procedure (2). I n the following paragraphs the d(3t:iils and limiting conditions for the high frequency titration arc ~ I Y srnted. TITRATION CONDITIONS
Reagents. Mercuric nitrate (0.01 M ) was prepared by dissulving the reagent grade salt in water with excess acid to give H. concentration between 0.004 and 0.006 M in nitric acid. Escess acid was required to prevent hydrolysis and precipit.ation of basic mercuric salts. Thi? mercuric nitrate solutioii was standardized as 0.01086 ,If with 0.01OOO M potassium thiocyanate prepwed from Acculute solution. Standard 0.01OOO M , sodium chloride was also prepared from Acciilute solution and cbecked against the thiocyanate solution by titrating both against equal portions of a 0.03 1 ' .i silver nitrate solution. All concentrations checked within 0.1 %. Titration Procedure. I n all the following titrations, various aliquots of the standard sodium chloride solution were titrated with the standard mercuric.nitrate solution. The titrations t,ook place a t room temperature (20' to 25' C.), and the end point, volumes fell between 70 and 80 ml. The differential titration procedure using the 30-mc. titrimeter was followed in all titrations (6). All titration curves obtained were similar to curve C, Figure 1, although in some cases where larger amounts of acid or interferences were present, the magnitude of the change at the end point was reduced. Results of all titrations performed arr summarized in Table I.
However, consideration of the sthichiometry of this renctioii also showed that it caused no stoichiometrical error in thc mercurimetric titration of chloride. Within experimental error, this There are limitations to the use of the high frequriicv titiimwas verified by titrating 20 ml. of standard 0.01 M sodium chloeter; the principal one is that titrations are restricted to a ride with a standard solution containing 0.00920 M mercuric certain concentration range for adequate sensitivity ( 6 ) . For the perchlorate and 0.005 M perchloric acid to prevent hydrolysis and 30-mc. titrimeter, this range corresponds to 0.003 to 0.02 d l precipitation of basic mercuric salts. The end point volume was sodium chloride or 0.0007 to 0.004 If hvdrochloric acid. All dcabout 80 ml. The end point for these dilute solutions. as shown in curve A . Figure 1, was not good. I n addition, Table 1. Etrect nf Varying Conditions on End Point in Mercurimetric the potential did not reach equilibrium Determination of Chloride rapidly in the region of the end point. In this titration, the caloinel half-ccll Vol. 0.01000 .li " 0 8 Concn. a t HgCI? Concn. a t vel. 0.01086 M HdNO:)? Used& Detn. S a C l Used, MI. E n d Point, ,If E n d Point, M Observed Theoretical Error was iso1,ated from the titrated solution 0.0004 2.30 1 500 0.0002 2.30 0 ,00 by means of a bridge containing 0.1 M 2 0.0004 2.27 5 00 0.0015 2.30 -0.03 pcrchloric acid t o prevent contamination 0.0004 2.30 2.33 5 00 0.0030 f0.03 0.0016 0.00004 0.46 0.47 4 1 00 +0.01 of thc solution by cliloricle. The poten0.0016 5 5 00 c .COO2 2.30 2.30 0.00 0.0016 2.31 6 5 00 +0.01 0.0015 2.30 tiometric titration is apparently not a 0.0016 5 00 0.0030 2.30 2.33 7 +0.03 good one. 0.0032 5 00 2.31 0.0015 2.30 8 +0.01 0.0064 2.25 9 2 00 O.OO15 2.30 -0.05 The mercurimetric titration of chlo0.020 0.0015 5 00 2.30 10 ?a ?" 0.000s 4.45 11 10 00 0.0004 -0.04 4.60 ride does not, appear t o have been per9.25 0.0016 12 20 00 9.20 0.0007 4-0.05 'formed conductometrically, although 0.0020 11.52 11.51 13 25 00 0,0009 fO.01 0,0040 22.97 23.01 0.0018 14 50 00 -0.04 this is theoretically possible. A conductometric titration was performed Indefinite end point. undor the same conditions as described ~~
0
-
1412
terminations in Table I fell within this range, except 1 and 4, where potassium nitrate was added us an inert electrolyte to make the sensitivity adequate. Effect of Acidity on the End Point. Within the range of conditions which could be studied, acidity has a negligible effect on the end point, as shown in determinations 1 to 3. ,This is not the case where chemical indicators are used to establish the end point. Acidities lower than that, in determination 1 could not be studied conveniently, for this acidity resulted from, the nitric acid introduced with the standard mercuric nitrate solution. Higher acidities were obtainqd by adding required amounts of 0.0100 .If nitric acid to the samples before titration. In these determinations, the mercuric chloride present a t the end point corresponded to that formed in titrating the sodium chloride sample; no mercuric. chloride was added. It was not possible to study directly the effects of low acidity for larger sample sizes because of the appreciable amounts of acid added with the standard mercuric nitrate solution. However, end point concentrations of mercuric chloride corresponding to larger samples of sodium chloride-Le., 20 ml.-were achieved by adding reagent, grade mercuric chloride as 0.100M solution in determinations 4 to 7. Acidities corresponding to much more than 0.003 I\. nitric acid could not, be studied, for these exceeded the optimum concentration range for good sensitivity of the titrimeter (2). It is apparent that the end point, acidity may lie anywhere between 0.00004 M and 0.003 M without causing error in this titration. The upper limit is set by inability of the 30-mc. titrimeter to handle higher concentrations. When a titrimeter of higher frequency becomes available, larger aridities may be tolerable. Below O.ooOo4 M , there is an indication, obtained by titration of samples even smaller than that in determination 4,that results are high-Le., too large a volume of mercuric nitrate is required. This is probably due to hydrolysis of mercuric ion a t such low acidity. Effect of Mercuric Chloride Concentration on End Point. When indicators are used to est,ablish the end point, the sample size has an effect on the size of the blank, owing to the formation of mercurous chloride ions (1, 3). I n the high frequency titration, this causes no discrepancy between the observed and theoretical end points, as shown by determinations 2, 6, and 8 to 10. However,, the end p i n t inflection was spread out increasingly,. until, for determination 10, the end point could no longer be located precisely. (This does not detract from the method; the mercuric chloride concentration of determination 10 would never be achieved by titrat,ing with 0.01 A I mercuric nitrate.) Effect of Sample Size’on the End Point. To check the conclusions of the two preceding sections, different sized samples of sodium chloride standard solution were titrated with the mercuric nitrate, as shown in determinations 1 and 11 to 14. KO excess nitric acid or mercuric chloride was added; the concentrations of these two substances corresponded to the concentrations added with the shndard solutions and the concentrations formed in the titration. These data confirm the agreement between end and equivalence points over a wide range of conditions. Interferences. Theoretically, any catibn may interfere which complexes or precipitates chloride ion with the same order of efficiency as mercuric ion. Also, any anion may interfere which complexes or precipitates mercuric ion with the same order of efficiency as chloride ion. The possible interferences of several kinds of ions were studied by adding them in roughly equivalent quantities to identical 25ml. aliquots of a 0.01000 M sodium chloride! solution and titrating these aliquots with standard 0.01086 U mercuric nitrate. Results are shown in Table 11. The anions studied were phosphate, acetate, oxalate, borate, fluoride, and cyanide. Of these, none caused appreciable error in acidic solutions except cyanide. In cyanide a ’single, sharp end point was observed, corresponding approximately to the sum of chloride and ,cyanide. However, the sharpness of the end point was slightly decreased by all these anions except phosphate. Iodide, bromide, and thiocyanate are known to bind mercuric ion even more effectivelythan chloride; these were not studied. The cations studied were cupric, ferric, and plumbous. Table 11 shows that none of ‘th‘ese cations interfered under conditions of these experiments. Silver ion was not studied as it is known to bind chloride even more effectively than mercuric ion. This study of interferences is not complete, but it is indicative. As the ratio of interfering ions to chloride increasesj the error undoubtedly becomes appreciable. This is probably particularly
ANALYTICAL CHEMISTRY true in lead, which forms an undissociated chloride, and also in oxalate, which precipitates mercuric ion. The limiting ratios for negligible interferences were not studied; these would probably depend on sample size and acidity. In any intended application of the method, all potential interferences should be studied before use of the method. There are usually too many considerations involved to predict safely whether or not a particular substance will interfere.
Table 11. Interferences in t h e Mercurimetric D e t e r m i n a t i o n of Chloride [In all determinations 25.00 ml. of 0.01000N NaCl containing the interfering substance was titrated with Oh1086 M Hg(NOa)z] Moles Interfering Ion Interfering Titration, 34ode of Ion/Mole M1. Standard Type addition Chloride Hg(NOdz Reqd. None .... None 11.50 Phosphate 11.53 0.5 Hap04 1.0 HOAc Acetate 11.49 0.5 Oxalate HzCzO4 11.48 Borate 1.0 HaBO: 11.53 KF Fluoride 2.0 11.50 KCN 1 .o Cyanide 22.3 cupric 0.5 11.53 Cu(N0a)r Ferric 0.5 11.50 Fe(N0da Plumbous 0.5 11.48 Pb(N0dz
CONCLUSIONS
Table I shows it is possible to determine chloride directly by high frequency titration with a precision corresponding to 0.03 ml. of 0.01 M mercuric nitrate. There is no appreciable bias on the results. Accurate results are obtained over a wide variety of conditions when the 36-mc. high frequency titrimeter is used. Xitric acid concentration may range from 0.00004 to 0.003 M a t the end point. Sample size has no effect on accuracy, as long as the mercuric chloride concentration formed a t the end point does not exceed 0.006 M ; above this concentration, the sharpness of the end point decreases. KO variable blank corrections are needed. Interferences do not appear to be many. For the mercurimetric titration of chloride, the high frequency titrimeter appears to be superior to the conductometric or potentiometric methods for establishing the end point. This procedure illustrates how the high frekuency titrimeter may be used to simplify existing titration procedures. It must be stated unequivocally that the mercurimetric determination of chloride is inferior to the argentimetric determination, providing a precipitate is not objectionable, since the end point is considerably sharper for the latter procedure (9). The effect of temperature on the end point could not be systematically studied without inconvenience. However, the end point is not critically dependent on temperature, for the titrations in Table I were carried out a t room temperature, which varied from 20’ to 25 C. ACKNOWLEDGMENT
This work was supported in part by the Wisconsin Alumni Researrh Foundation. LITERATURE CITED
(1) Blaedel, W. J., and Lewis, W. B., unpublished data, Nurth-
western Uqiversity, Evanston, Ill., 1942.
(2) Blaedal, W. J., and Malmstadt, H. V., AN.AL.CHEM.,22, 734 (1950). (3)
Kolthoff, I. M., and Sandelt, E; B., “Textbook of Quantitative Inorganic,AnrtlysiJ,”pp. 479, 576, New York, Macmillan Co., 1946.
Latimer, W. M., “Oxidation Potentials,” pp. 162-7, 177, New York, Prentice-Hall, Inc., 1938. (5) Miiller, E., and Aarflot, H., Rec. t r m . ehim., 43, 574 (1924). (6) Roberts, I., IND.ENO.CHEM.,ANAL.ED.. 8, 365 (1936). (4)
RECEIVED March 20, 1950.
.
