743
V O L U M E 26, NO. 4, A P R I L 1 9 5 4 However, a decrease in titer of 1% in four months on storage in soft glass bottles has been reported by Goeta, Loomis, and Diehl (9).
Bets, J. D., and Noll, C. A , , J . Bm. Water Works Assoc., 42, 49,749 (1950).
Biedermann, W., and Schwaraenbach, G., Chintia (Switz.), 2, 56 (1948).
TITRATIONS WITH PRIMARY STANDARDS
Blaedel, W. J., and Knight, H. T., ANAL.CHEY.,26,743 (1954). Cheng, K. L., Kurta, T., and Bray, R. H., Ibid., 24, 1640 (19.52). ,----,-
Solutions of 0.01F NazHzYwere prepared from the dihydrate and anhydrous disodium salt as primary standards, and were titrated against standard solutions of copper nitrate, calcium chloride, and zinc chloride, using the high frequency technique to avoid indicator errors and corrections. These titrations are described in another paper (4). The results of 20 titrations gave an assay value of 100.00% Na2H2Y for the primary standard anhydrous form, and five titrations gave an assay value of 100.05% Na,HZY. 2H2O for the primary standard dihydrate. The relative standard deviation of a single titration was around 0.1%.
Diehl, H., Goetz, C. A., and Hach, C. C., J . Am. Water Works Assoc., 42,40 (1950). Flaschka, H., Mikrochemie w r . Mibrochim. Acta, 39, 38 (1952). Ibid., p. 315. Goetz, C. A., Loomis, T. C., and Diehl, H., ANAL.CHEM.,22,
ACKNOWLEDGMENT
Pribil, R., and Koudela, Z., Collection Czechoslou. Chem. Corn-
The authors wish to thank E. I. du Pont de Nemours B: Co. for a grant-in-aid, and the Atomic Energy Commission for a research grant in support of this work.
Pribil, R., and Matyska, B., Ibid., 16, 139 (1951). Schwarsenbach, G., and Biedermann, W.3 Helv. Chim. Acta,
798 (1950).
Harris, mi. F., and Sweet, T. R., Ibid., 24,1062 (1952). Hernandez, H. R., Biermacher, U., and Mattocks, A. JI., Bull. Natl. Formularv Comm., 18, 145-52 (1950). Keihei, U., ANAL.CHEM.,24, 1363 (1952). Langford, K. E., Electroplating, 5 , 41 (1952). Mattocks, A. M., and Hernandez, H. R., J . Am. Assoc. Pharrn. (Sci. Ed.). 39.519 (1950). h l k g e r , J.’R., kippier, R’. W., and Ingols, R. S., rlx.4~.CHEX., 22,1455 (1950). mum., 16,80 (1951).
31,459 (1948).
Schwarsenbach, G., Biedermann, W., and Bangerter, F., Ibid., 29,811 (1946).
LITERATURE CITED
(1) Ranewicz, J. J., and Kenner, C. T., ASAL. CHEM.,24, 1186 (1952).
Weissberger, A,, “Physical Methods of Organic Analysis,’* New York, Interscience Publishers, 1949. RECEIVED for review July 13, 1953. Accepted December 2, 1953.
Stoichiometry of Titration of Metal Ions with Disodium Salt of Ethylenediaminetetraacetic Acid Using High Frequency Technique W. J. BLAEDEL and H. T. KNIGHT‘ Department o f Chemistry, University o f Wisconsin, Madison, Wis.
I
T IS the purpose of this study t o show that direct titrations of some typical metal ions [copper(II), zinc, calcium, and magnesium] with the disodium salt of ethylenediaminetetraacetic acid (NazH2Y) are stoichiometrical over rather wide ranges of condit ons, within relative errors of 0.1%. The titration of many kinds of metal ions with the disodium salt of ethylenedianiinetetraacetic acid is described in the literature (1) usually a t error levels of 0.5 to 2%, though in some cases greater accuracy is claimed. Very recently, it has been shown that the disodium salt of ethylenediaminetetraacetic acid standardized against copper as a primary standard reacts stoichiometrically with nickel and iron, a t the 0.1% error level ( 5 ) . I n most of these works, the titrant is standardized by titration against known solutions of the sought-for metal, or of another primary standard metal. If properly performed, this method is acceptable for circumventing errors due to reagent impurities, indicator blanks, side reaction, and nonstoichiometrical reaction between the reagent and sought-for substance. However, this method does not easily reveal the magnitudes and sources of the errors, nor does it allow the conclusion that the reaction on which the titration is based is stoichiometrical. In proving the stoichiometry of any titration, an instrumental method of following the course of the reaction is preferable to the use of chemical indicators. Elimination of the indicator error allows easier resolution of the other errors mentioned above. In the particular case of the direct titration of the alkaline earths in ammoniacal solution, the differential high frequency titration method (.2) is very advantageous, since the endpoint break is quite 1
Present address, Fansteel Metallurgical Corp., North Chicago, Ill.
indistinct for the ordinary type of conductometric titration curve obtained by plotting conductance against volume of standard solution. TITRATION OF COPPER
-40.0094331 solution of the disodium salt of ethylenediaminetetraacetic acid was prepared from the dihydrate as a primary standard (1). A standard solution of copper nitrate was prepared from copper foil (assay value, 99.98% copper). The foil was cleaned first with ether, then with dilute nitric acid, then rinsed off with distilled water and air-dried before weighing. The weighed foil was dissolved in 5 ml. of redistilled water and 2.5 ml. of concentrated nitric acid and diluted to 1 liter in a volumetric flask, giving a solution which was 0.01000M in copper nitrate and about 0.02.V in nitric acid. Titrations of aliquots of the standard copper nitrate with the standard disodium salt were performed in end-point volumes of 80 to 100 ml., since the titration cell required 50 to 100 ml. of solution. Also, the concentration range for adequate sensitivity of the 30-megacycle instrument lies between 0.0001 and 0.01M sodium chloride, or solutions of other electrolytes with the same conductances. It was for this reason that the investigation was confined to titrations with 0.01M, rather than 0.1M solutions. The titration technique was quite similar to that described in an earlier publication ( 9 ) . In the region of the equivalence point, standard disodium salt was added a t a uniform rate (about 0.5 ml. per minute), and the rate of change of frequency was followrd on a recording milliammeter, giving a recorded differential titration curve. Results of titrations of different-sized aliquots of standard copper nitrate a t different acidities are shown in Table I. Differential curves for some of the titrations of Table I are shown in Figure 1. In Figure 1, abscissas represent volume of standard disodium salt added, and ordinates represent a current which is proportional to the rate of change of conductance produced by
744
ANALYTICAL CHEMISTRY
__----
15.7
-I
16.1
-\
B
c
-1 21.0
I/+---
__----
‘21.4
D
of a single titration corresponds to 0.016 ml. of 0.01M Ka2H2Y. The titration of copper(I1) is therefore stoichiometrical within the precision of the work. 2. The titration of copper(I1) may be performed with good accuracy down to pH 2.5. At p H below 2.5, the end point becomes diffuse,owing to the decreased stability of the CUT-- complex. Also, the magnitude of the high frequency end point change is decreased-Le., the sensitivity is decreased owing to the high conductance of the acid solution. 3. The titration of copper(I1) may be performed with high accuracy in ammoniacal solution. The use of concentrated ammonia is particularly advantageous for the high frequency method, since ammonia does not contribute appreciably to the conductance of the solution. The ammonia may be as concentrated as 2M TTithout interference, so long as enough is present to prevent precipitation of cupric hydroxide. End points cannot be successfully established by the differential technique when copper(I1) is in the form of a precipitate, apparently because the precipitate does not dissolve rapidly in the region of the end point. The ammonia used as a reagent in this work was redistilled and stored in a polyethylene bottle. When ordinary reagent grade concentrated ammonia was used, there was a blank corresponding to about 0.1 ml. of 0.01M Na2H2Yper 10 ml. of concentrated ammonia.
Table 11. Titration of -4liquots of 0.0100M Zinc Chloride Solutiona w-ith 0.00943M Na2HSY
1.00
Other Reagents Addedb 1 ml. redist.
4.98
1 ml.. redist.
ZnCIz, MI.
End Point pH
Sa2HzY
3.5
21.17
Observed 1.07 1.06 5.31 5.32 10.60 10.60 15.89 15.89 21.18 21.20 26.47 26.48 21.15
2 5
21 17
Obscure
10
”2
ML. STANDARD
Na2H,Y
9.98
Figure 1. Some Differential Curves for Titration of Aliquots of 0.01000M Copper Nitrate with the 0.00943M Disodium Salt of Ethylenediaminetetraacetic Acid
14.99 19.97 24.99
For legend, see Table I
adding the disodium salt solution a t a uniform rate. A pair of reference volumes is given for each curve, and the dotted ordinate representing a zero rate of frequency change is also shown. The following conclusions may be drawn from the work associated with Table I and Figure 1. 1. From the 15 titrations in Table I, the median error corresponds to 0.00 ml. of 0.01N Na2H2Y,and the standard deviation
Table I. Titration of Aliquots of 0.01OOOM Cbpper Nitrate Solution“ with 0.00943M Na2H2Y cu(XO~)z,Other Reagents hI1. Addedb 1.OO 1 ml. redist. NH3
NaaHaY Required, hll. TheoOhretioal served 1.06 1.06
4.98
1 ml.
redist.
10
5.28
9.98
1 ml.
redist.
10
10.58
14.99
1 ml.
15.89
1
redist. XHa ml. redist.
10
19.97
10
21.17
10
42.34
h*Ha
SHs
“8
39.94 19.97 19.97 19.97
1 ml.
redist. h’H3 10 ml. redist.
SH3 20ml.0.02.V NaOH
0.00 0.00 0.01 0.01 -0.01 0.01 0.00 -0.02 -0.03 0.01 -0.04 0.00
0.03
21.17
21.20
3.5
21.17
21.17
2.5
21.17
21.16
11
Error, hll.
5,29 5.29 10.57 10.59 15.89 15.87 21.14 21.18 21.13 42.34
1.06
Curve in Fig. 1
0.00 -0.01
Also contains about 0.02.M HNOa. b Water added to make end-point volume 80 to 100 ml. in all cases.
5
PiH3 1 ml. redist. SH3 1 mi. redist. SH3 1 ml. redist.
10
5,28
10
10.58
10
15.89
10
21.17
10
26.50
“3
19,97
End Point pH 10
SH3 1 ml. redist.
A
B C
D E
19 97
20 ml. 0.020.M SaOAc
Required,
Theoretical 1.06
Error, hll. 0.01 0.00 0.03 0.04 0.02 0.02 0.00
0.00 0.01 0.03 -0.03 -0.02 -0.02
Also contains about 0.004.M HC1. b Water added to make end-point volume 80 t o 100 ml. in all cases.
5
TITRATION OF ZINC
A standard zinc chloride solution x a s prepared from zinc oxide as a primary standard ( 4 ) . Reagent grade zinc oxide was ignited for 2 hours a t red heat over a Weker burner. After cooling in a desiccator over calcium chloride, a 0.8138-gram portion was weighed out, dissolved in 5 ml. of water plus 2.0 ml. of concentrated hydrochloric acid, and made up to 1 liter in a volumetric flask, giving a solution which \vas 0.01000M in zinc chloride and about 0.004M in hydrochloric acid. Aliquots of the standard zinc chloride were titrated with the same standard disodium salt used for the previously described titrations of copper nitrate, and using the differential high frequency method for establishing the end point. Results of titrations on different sized aliquots of the standard zinc chloride solution at different acidities are shown in Table 11. The differential titration curves from which these end points were taken were very similar to those for copper (Figure l), and are not reproduced in this paper. The following conclusions may be drawn from the work asso-, ciated with Table 11. 1. From the 13 titrations in Table 11, the median error corresponds to 0.01 m]. of 0.01U Sa2H2Y,and the standard deviation of a single titration corresponds to 0.020 ml. of 0.01N Sa2H2Y. The titration of zinc (Zn ++) is therefore stoichiometrical, within the precision of this work. 2. The titration of zinc(I1) may be performed with good accuracy down to pH 3.5. A4tpH beloF 3.5, the end point be-
V O L U M E 26, NO. 4, A P R I L 1 9 5 4
745
comes diffuse, owing to the decreased stability of the ZnT-complex. 3. The titration of zinc(I1) may be performed with high accuracy in ammoniacal solution, with the same advantages as described for the titration of copper(I1). However, a large excess of ammonia should be avoided: the end point becomes less sharp for ammonia concentrations exceeding 0.5M, probably owing to formation of Zn(?l”s)d in competition with formation of ZnT-in the region of the end point. A satisfactory procedure is to add enough concentrated ammonia to dissolve any zinc hydroxide which first forma, and then to add about 1 ml. in excess. Redistilled ammonia should lie used, or a correction should be applied if ordinary reagent grade concentrated ammonia is used.
hydrochloric acid was back-titrated with standard barium hydroxide. ] Also, duplicate analyses of the same lot by measurement of the weight loss on ignition to calcium oxide gave 99.98 and 99.94% calcium carbonate. Since making these tests, another lot of S.L. grade calcium carbonate mas used to prepare mixtures of calcium carbonate and calcium sulfate, which have been used as unknomms by students to measure weight loss on ignition. Results over two semesters (140 students) indicate an assay value over 99.95% calcium carbonate. These results on only two lots of S.L. grade calcium carbonate can only indicate the adequacy of this material as a general primary standard for calcium, but it seems reasonable to conclude that the particular lots used are suitable primary standards for testing the stoichiometry of the reaction between calcium and the dipodium salt of ethvlenediamine tetraacetic acid.
TITR.4TION OF CALCIUM
Table 111. Titration of Aliquots of 0.01000M Calcium Chloride Solution” with Standard R’a?H?YSolution N.?.FT.Y - .-.- - .-
A standard calcium chloride solution was prepared from a special grade of calcium carbonate as a primary standard. Reagent grade calcium carbonate, and reagent grade calcium carbonate labeled “low in alkalies” could not be reliably used as primary standards for calcium. .4nalysis of various lots acidimetrically, gravimetrically, and by titration with Sa2H?T indicated discrepancies of 0.1 to 0.5y0in the calcium carbonate assay values. Homvcr, Mallinckrodt standard luminescent (S.L.) grade calcium carbonate is manufactured to meet very high purity standards. IVhen a sample of S.L. grade calcium carbonate was dried for 2 hours at 130” C., an acidimetric assay gave 99.95% calcium carbonate. [The calcium carbonate was dissolved in excess Ptandard hydrochloric acid and the excess
Other Reagents Addedb 1.0ml.redist. SHa
End Point pH 10
9.98
1.0ml.redist.NH1
10
14 99
l.Oinl.redist.SH3
10
10.97
1.0rnl.redist.SH~
10
CaCh, 311.
1.00 4.98
24 99
1.0 nil. redist. SHa
19.97
0.25
..
nil.
redist.
10 ,
.
Required, 111. ObTheoserved retical 1.06C 1.07 5.28c 5.29 5.31 10.58C 10.59 10.60 15.8gc 15.89 15.89 2 1 . 1 7 c 21.19 21.18 21.19 26.5OC 26.50 26.53 21.17C 21.19
..
21.17C
19.9i
\-U“ __” 10 nil. ethylenedianiine 6 0 1111. 0 1031
10
19.67d
19.97
S a~.~~ OH 9 0 nil. 0.1O.V
11
19.97e
19.97
Error, 111.
0.01 0.01 0.03 0.01 0.02 0.00 0.00 0.02
o.oi
Curve in Fig. 2
A B C
0.02 0.00 0.03 0.02
21.16 21.17 19.64
-0.01 0.00 -0.03
19.95 19.94
-0.02 -0.03 -0.01
D
~
NaOH 1 0 . 0 ml. 0.10.V 11 19.97e 1 9 9 6 KaOH 20.01 0 04 Also contains about 0.004M HC1. b Water added t o bring end-point volume to 80 to 100 nil. 0.00943.M NazH2Y. d 0.01015M KazHIY. e 0.01OOOM NasH2Y.
19.97
E
Q
-
-
-
____ _ _
21.0
21.4
1-
1
19.4
----_---
-1
19.9
19.7
20.2
~~
ML. STANDARD Na2H,Y
Figure 2. Differential Curves for Titration of 0.01000M Calcium Chloride with Standard Disodium Salt of EthylenediaminetetraaceticAcid For legend,
see
Table I11
A standard calcium chloride solution was prepared by dissolving an accurately weighed portion of the dried S.L. grade of calcium carbonate in 5 ml. of water plus 2.0 ml. of concentrated hydrochloric acid and diluting to 1 liter in a volumetric flask, giving a solution containing 0.01OOON calcium chloride and an excess of about 0.004M hydrochloric acid. Various sized aliquots of this calcium chloride solution were titrated under various conditions against standard disodium salt solutions. Most of the titrations were made \
