Determination of Ammonia with Cupric Carbonate R. C. BLlNN and F. A. GUNTHER University of California Citrus Experiment Station, Riverside, Calif.
b Two methods for the determination of ammonia based upon the solubility of ammonium salts in alkaline solution are presented: a titrimetric procedure in which precipitated cupric carbonate is completely dissolved by titration with a solution of the ammonia sample, and a colorimetric procedure in which a solution of the ammonia sample is added to excess precipitated cupric carbonate and the resultant dissolved blue cupric ammonium salt is determined (Anlax. 700 mp).
R
investigations have shown that gaseous ammonia from ammonia-releasing formulations is promising in control of the growth of the blue-green molds PeniciUium italicurn Wehmer and P. digitaturn Saccardo on citrus fruits during shipment and storage (1). These formulations consist of an ammonium salt admixed with an alkaline material, along with excipient material such as bentonite and talc as lubricating, bulking, and moisture-attracting aids. To devise the most efficient formulation for the proper release of ammonia, analytical procedures for total ammonium salts were required. The usual procedure involves a Kjeldahl-type distillation from alkaline solution for acidimetric titration. Two new methods based on the solubility of copper ammonium salts in alkaline solutions are presented here. ECEKT
TITRIMETRIC PROCEDURE
All reagents are analytical reagent grade. CUPRIC SULFATE. Dissolve 3.3 grams of cupric sulfate pentahydrate in 100 ml. of water. Standardize against ammonium sulfate solution before use. AMMONIUM SULFATE. Dissolve 1.50 grams of anhydrous ammonium sulfate in 100 ml. of water to use for standardizing the cupric sulfate solution. Other concentrations should be used for establishing a standard titration curve. DILUTESULFURIC ACID. A concentration should be used that will result in a slightly acidic solution of the ammonium compound to be tested, 0.2iM in the present study. Procedure. hiix the ammoniacontaining sample thoroughly with a measured portion of the dilute sulfuric acid solution until all the ammonium Reagents.
1882
ANALYTICAL CHEMISTRY
salts are in solution, then gravityfilter the extract through filter paper. The extract is stable for storage under these conditions. Immediately prior t o titration, add an excess of anhydrous sodium carbonate to this acid solution and filter through cotton into a 5- or 10-ml. buret. In the meantime, mix 1 ml. of 10% sodium carbonate solution with an exact, predetermined amount of standardized cupric sulfate solution (0.25, 0.5, 1.00 ml., etc.). To this mixture add dropwise the clear ammonia solution until a permanent clear-blue solution results. The end point may be accentuated by use of a pin point of light shining through the solution (2). The amount of ammonia present is determined from a standard titration curve (Figure 1) prepared from various concentrations of standard ammonium sulfate solution. COLORIMETRIC PROCEDURE
Reagents. CUPRICSULFATE.Dissolve 10 grams of cupric sulfate pentahydrate in 100 ml. of water. CALCIUM NITRATE. Dissolve 22 grams of calcium nitrate tetrahydrate in 100 ml. of water. Procedure. Transfer in sequence into a 50-ml. glass-stoppered Erlenmeyer flask 10 ml. of 4% sodium carbonate solution, 2 ml. of cupric sulfate solution, 1.0 ml. of the ammonium salt sample solution, and 1 ml. of calcium nitrate solution, mixing thoroughly after each addition; then filter. Determine the absorbance of the clear filtrate a t 700 mp with any type of spectrophotometer or compare it visually with previously prepared color standards. A standard curve may be prepared from a standard ammonium sulfate solution (Figure 2). The colored solution is stable for more
than 24 hours if tightly stoppered to prevent loss of gaseous ammonia. DISCUSSION
Titrimetric Procedure. This method depends on the complete solution of precipitated cupric carbonate as the cupric ammonium salt by titration with the ammonia sample. A standard titration curve for 67 mg. to 667 mg. of ammonia dissolved in 100 ml. of dilute acid solution is presented in Figure 1. This range can be extended by using different quantities of cupric carbonate or by varying the amount of hydrochloric acid solution used to dissolve the sample. As can be seen from Figure 1, a few milliliters of sample extract would be sufficient to obtain a replicated titration. Small amounts of material may be assayed easily by adjusting
0.3 0.3
I
0.2
!
1
1
Figure 2. Standard calibration curve for colorimetric procedure
0.6-
E
0.5-
0
0 L
B
0.4 -
.g0 0 . 3 E
5 0.2 m
u
0.1
-
0.0 1 0.0
Figure 1. samples
I .0
1
0.4 0.6 0.8 1.0 1.2 Grams Ammania per 100ml.
2.0
I 3.0
Titration, mi. Standard titration curve, based on 100-mi.
the dilutions and the amounts of cupric carbonate. The addition of excess sodium carbonate before final assay is necessary where incidental materials in the formulation sample-such as calcium chloride-would precipitate as carbonates, obscuring observations of the solution of the copper ammonium salt. Interfering compounds are copper salts which do not readily form copper ammonium complexes, and metals which do form ammonium complexes, including nickel(II), zinc(II), silver, and cadmium(I1). The end point, which is sharp, is easily detected with a little experience, and is distinguished by the sparkling clearness of the solution. Duplicate titrations are recommended. Table I lists duplicated assays of a typical ammonia - releasing formulation containing ammonium sulfate, sodium carbonate, bentonite, and talc. Colorimetric Procedure. This method is based upon the solubility of insoluble cupric carbonate as the blue cupric ammonium complex in alkaline solution. Because cupric carbonate is partially soluble in the presence of excess bicarbonate and carbonate ions, calcium nitrate is used t o precipitate excess carbonates. A standard curve prepared from standard ammonium sulfate solution ranging from 0.2 to 1.2 grams of ammonia dissolved in 100 ml. of water is presented in Figure 2. While this curve does not follow Beer's lam, the
Table 1.
Assays of a Typical AmmoniaReleasing Formulation"
cuso, Soln., M1. 0.50 1.00 0.50
Titrant, M1. 0.55 1.35 0.55 1.25 0.60 1.35
Total "a,
Mg. 35 33 35 1.00 35 0.50 33 1.00 33 Av. 34 Winn-Mat fungicide KO.3 (WinningPeplow, Inc., Los Angeles, Calif.).
Table 111. Determination of Ammonia in Synthetic Ammonia-Releasing Formulations"
KH1 Added, Gram 0.20
NH1 Found, Gram Titrimetric Colorimetric 0.21 0.20 0.21 0.21 0.21 0.20 0.40 0.41 0.42 0.41 0.40 0.41 0.41 0.60 0.59 0.60 0.59 0.62 0.59 0.60 a Also contained sodium carbonate, bentonite, and talc.
Table II. Assays of a Typical Ammonia-Releasing Formulation'
Absorbance at Available NH, 700 mp Gram 0.575 0.610 0.565 0.600 0.605 0.570 0.613 0.580 0.613 0.580 0.613 0.580 0.623 0.590 *' Formula tion WB-10 (WinningPeplow, Inc., Los Angeles, Calif.).
slight deviation from a straight line does not interfere with the accuracy of the method. The sensitivity range can be extended by using different quantities of reagents. This colorimetric procedure is subject to the same interferences as the titrimetric proEedure. Table I1 lists replicate assays
of another typical ammonia-releasing formulation containing ammonium carbonate, bentonite, and talc. The recoveries of ammonia from synthetic formulations by the two procedures are compared in Table 111, indicating satisfactory recovery. LITERATURE CITED
Gunther, F. A., Kolbezen, hl . Blinn, R. C., Staggs, E. Barkley, J. H., Wacker, G. Klotz, I. J.. Roistacher, C. El-Ani, A., Phytopathology 46, (1956). Kolbezen, M. J., Staggs, E. Drivate communication. RECEIVEDfor review March 15, 1957. Accepted August 12, 1957. Division of Analytical Chemistry, 131st Meeting, A('S Miami, Fla., April 1957. Paper 970, Cniversity of California Citrus Experirnent Station, Riverside, Calif.
Determination of Small Amounts of Iodide in the Presence of Chloride by Potentiometric Titration R. H. STOKES and 1 A. WOOLF Deparfmenf o f chemistry, University o f New England, Armidale, N.S. W., Ausfralia ,The iodide end point in the potentiometric titration of iodide-chloride mixtures with silver nitrate can readily be determined in the presence of a large excess of chloride, though small amounts of chloride interfere. This is explained in terms of the effect of the formation of complex argentochloride ions in delaying the precipitation of silver chloride.
T
investigation arose from the need to determine small quantities of iodide in the presence of large concentrations of chloride in connection with tracer-diffusion studies of iodide ion, to be reported elsewhere. HIS
Potentiometric titration of 0.01M iodide with silver nitrate was perfectly satisfactory, if no chloride a t all, or a large excess of chloride-e.g., 1Mwas present. Smaller concentrations of chloride-0.1M-caused serious interference, the iodide end point being ill defined, or in some cases well defined but several per cent from the correct value. This difficulty was overcome by adding a large excess of chloride before titration, in cases where the original chloride concentration was too low. This article gives a theoretical explanation of this rather surprising situation. EFFECTS OF COMPLEX IONS
To obtain a well-defined end point in a
potentiometric titration it is desirable to proceed several per cent past the equivalence point, and then plot AE/Az ( E = e.m.f., z = volume of titrant) us. x,obtaining a peak a t the equivalence point. This peak will not be well defined if precipitation of silver chloride begins before or too soon after the iodide equivalence point. The calculations which follow show that the precipitation of silver chloride is actually delayed by the presence of a large excess of chloride in the solution, owing to the existence of complex argentochloride ions. The situation is thus very different from what would be expected on the simple basis of comparing the solubility products of silver chloride and silver iodide. VOL. 29, NO. 12, DECEMBER 1957
1883
