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 t h e 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 t h e solubility products of silver chloride and silver iodide. VOL. 29, NO. 12, DECEMBER 1957
1883
Silver chloride is appreciably soluble in concentrated chloride solutions and silver iodide even more so in concentrated iodide solutions. I n the present problem complexes of the type AgI, will be unimportant, as the iodide ion concentration near the iodide end point must be very low, and those of the type AgIC1- can be neglected in view of the observed low solubility of silver iodide in chloride solutions. Those of the type AgClp- are, however, very important. From the solubility data for silver chloride in alkali chloride solutions (S), it is easily shown that the equilibrium constant, K z , for the proc2C1- = AgC12- is of the ess Ag+ order of K 2 = 106 mole+ liter2. (Higher complexes occur, but do not materially alter the argument. The solubility of silver chloride in hydrochloric acid is of similar magnitude to that in alkali chloride solutions, so that p H is clearly not an important factor.) The other relevant constants are the solubility products, here denoted by
+
I
Now the e.m.f. of a silver-silver iodide electrode system in the solution is a linear function of In [I-], E
=
E"
+ k In [I-], (12
al
-x
sinh y
= 2 x/S'
-
= b, t
lhg'l
+ [a?$l;l
[AgCl;] = a:&
Also
[Ag+]
(2)
(3)
From Equations 1. 2, and 3 Ive obtain +
+ [I-] 1 +U~KZ
= x - a1
[Ag
(4)
Hence, because solid silver iodide is present,
+
Putting S1(l ai&) = S' for convenience, Equation 5 becomes: 1884
(7)
From Equation 8 n e see that the curve of E vs. x has a point of inflection a t the iodide equivalence point x = al. (This is the usual criterion for a potentiometric end point, found by plotting dE/dx us. x to obtain a peak a t the end point.) The above equations hold up to the moment when silver chloride precipitation begins. At this moment, as both solid silver chloride and solid silver iodide are present, =
[rlgL][I-] and
8 2
ANALYTICAL CHEMISTRY
1
0
O O l N KI 0 -
50
I N KCI
i I
=
uz [&+I
i.e., [I-] = S1a2/Sz (9) Substituting Equation 9 in 6 gives:
If the iodide end point is to be passed before silver chloride precipitation begins, Equation 10 must give positive values of x - a,. This will be the case, provided that
This condition is fulfilled for all values of a2 in the present case, since 8, = 10-l6, Sz = lO-"J, and K z = lo6. However, to obtain a well-defined iodide end point, silver chloride precipitation must not occur until the iodide end point is exceeded by a t least 1%, and preferably more-Le., we require: x-u1=
[I-]
(1) Since the total amount of silver in the system is L equivalents, present as iigI, Ag-, and AgCl;, we have T .
RT/F)
50
[&+I [CI-I
b, = a,
=
Therefore, putting g = In ([I-]/.\/S'), Equation 6 becomes :
81
(Throughout, subscript 1 denotes quantities involving iodide and 2 those involving chloride, IT-hile square brackets denote the molar concentrations of substances in solution.) Consider 1 liter of a solution containing initially a, moles of iodide and a2 moles of chloride, to which is added x equivalents of silver nitrate. The effects of activity coefficients will be neglected, as will the volume change occurring during the titration. It is assumed that solid silver iodide is present from a very early stage of the titration, but that silver chloride has not yet appeared, and conditions a t the iodide equivalence point are determined; these conditions are consistent with the assumed absence of 3. silver chloride precipitate. The amount of silver iodide precipitate (in moles) is
I
For values of a2 greater than about 0.01 mole liter-l, Equation 11 approximates well with the above value8 of 81, Sz, and Kz to:
$> 100 i.e., a large excess of chloride prevents precipitation of silver chloride near the iodide end point, as observed. Furthermore, the percentage by which the iodide end point can be exceeded is approximately proportional to the ratio of initial concentration of chloride to iodide. This result arises from the fortunate magnitudes of the equilibrium constants involved. For example, if the chloride complex did not exist, &-e should have
Figure 1. Effect of chloride on iodide end point
Kp = 0, and silver chloride would begin to precipitate a t the iodide equivalence point if az = X z / ~ / s l = 0.01;M and before it if u2 > 0.01V. SHARPNESS
OF IODIDE END POINTS
The readiness with which a potentiometric end point can be determined depends on the magnitude of the electromotive force change over a small region near the equivalence point, For ease of discussion \ r e shall consider the range from 1% before to 1% after the equiralence point-Le., 1: - a1 = -0.01 a1 to x - a1 = 0.01a1. Since by Equation 7 y = sinh-1-
a, - x 2dS'
(12)
the change in y (or in In [I-]) over this range is Ay = A In [I-] = 2 sinh-l
So that AE = ( 2 R T / F ) sinh-1 (200% ~
If a2 = 100 al, this becomes AE = (2RT/F) sinh-15
= (2RT/F)In 10 (13)
This is to be compared with the value in the absence of chloride, (at = 0), which gives AE = ( 2 R T I F )sinh-1
___
(20onts,,)
=((LETIF) In ( 106al)
which for al = 0.01 mole liter-’ gives AE‘ = ( 2 R T / F ) X 41n 10
Hence aE’/aE = 4, so that there is about a fourfold drop in the e.m.f. change near the end point as a result of the addition of a hundredfold excess of chloride ions. However, the change over the range *l% each side of end point is, by Equation 13, of the order of 100 mv., even when this large excess of chloride is present, so that the end point is easily determined. Figure 1 shows actual titration curves for three
equal samples of iodide, to one of which no chloride has been added, to another a hundredfold excess of chloride, and to the third a n equimolar amount of chloride. The erroneous result in the third case is readily apparent. These curves were obtained with a Muller ( 2 ) capillary electrode system, the actual electrodes being silver wires very lightly coated with silver iodide by anodic deposition from potassium iodide solution. The potentials were measured every 0.2 ml. near the end point by means of a potentiometer with a vacuum-tube voltmeter as null indicator. For routine determinations where actual voltages are unimportant, the simple battery-operated titrator described by Garman and Droz ( 1 ) is very satisfactory. Magnefic stirring is employed throughout the titration. The three types of titration show very
different visual appearances: In the absence of chloride, the silver iodide formed remains colloidally dispersed until it precipitates sharply a t the equivalence point; with a large excess of chloride the silver iodide remains colloidally dispersed throughout; and when only a small amount of chloride is present, the precipitate coagulates well before the equivalence point. LITERATURE CITED
(1) Garman, R. L., Droz, M. E., IND. ENG.CHEM.. ANAL. ED. 11, 398 (1939). (2) Muller, A., 2. physik Chem. 135, 102 (1928). (3) Seidell, A,, “Solubilities of Inorganic and Metal-Organic Compounds,” 3rd ed., p. 40, McGraw-Hill, New York, 1953.
RECEIVEDfor review May 14, 1956. Accepted July 8, 1957.
Flame Spectrophotometric Determination of Microgram Quantities of Magnesium H. STRUNK, and S. L. ADAMS Research Department, Joseph E. Seagram & Sons, Inc., louisville, Ky.
LOUIS MANNA, D.
F A rapid procedure is described for the flame spectrophotometric determination of micro quantities ( 1 to 6 y ) of magnesium. It is precise and requires no preliminary separations. The radiant power of magnesium was greatly enhanced by aspirating from an 80% acetone solution. The inhibitory action of aluminum was circumvented by the use of a multipleion radiation buffer consisting of 7 5 0 y of calcium per ml., 25 y of aluminum per mi., and 2M acetic acid. The accuracy of the method was verified by analyzing National Bureau of Standards samples of limestone and magnesite.
0
few papers dealing with the flame photometric determination of magnesium, only Dippel (6), Ikeda, (Q), and Knutson ( I S ) described procedures that can be applied to its microdetermination. Close, Smith, and Watson ( 1 ) determined the magnesium oxide content of Portland cement, limestone, and cement mortar after the separation of silica, iron, and alumina. In determining the magnesium content of glass, Roy (19) used synthetic standards corresponding to the known composition of the glass to be analyzed in order to overcome radiation interference. F THE
Kuemmel and Karl (14) similarly determined the magnesium content of cast iron. Pro and Xathers (18) used a radiation buffer of dextrose and phosphate in determining the magnesium content of wines. This study describes a rapid procedure for the determination of small concentrations (1 to 6 y per ml.) of magnesium. Although the method was developed principally to determine micro quantities of magnesium in beverage alcohol, it is applicable t o complex solutions containing many inorganic components without preliminary separations. APPARATUS AND CHEMICALS
Emission measurements were made with a Beckman Model DU spectrophotometer equipped with a Model 4300 photomultiplier accessory and a Model 9200 flame photometry attachment. A model 4020 atomizer-burner utilizing a n oxyhydrogen mixture was the source of excitation. Analytical grade chemicals were used throughout the investigation. Stock solutions of aluminum, copper, iron, and magnesium ions were prepared by reaction of each metal with a minimum of concentrated nitric acid and then diluting to volume with double distilled
water. Solutions of lead, cobalt, lithium, manganese, nickel, potassium, and sodium ions were prepared from their nitrates and a solution of calcium ions rvas prepared by reacting calcium carbonate with a minimum quantity of concentrated nitric acid. The solutions were stored in polyethylene containers. EXPERIMENTAL PROCEDURE
The 371- and 383-mp maxima in the magnesium oxide bands have been used for the determination of magnesium. However, unresolved radiation from the iron lines a t 373.0, 373.6, and 386.0 mp caused positive errors in the magnesium determinations. For this reason the much weaker atomic line a t 285.2 mb was used for the determination of this element. The following Beckman DU spectrophotometer settings were used throughout the investigation: Wavelength, mp Selector Sensitivity on photomultiplier battery box, 60 volts per dynode Resistor, megohms Slit width, mm. Hydrogen, lb./sq. inch Oxygen, lb./sq. inch
285.2 0.1
Full 22 0.025 1.25 12
The wave-length setting must be exVOL. 2 9 , NO. 12, DECEMBER 1957
* 1885
