1648
ANALYTICAL CHEMISTRY
very good, in view of the volatility of the compounds and the technical difficulties involved in sampling and in the analyses. I n no case did precipitate form in the trap solution of mercuric chloride (G in Figure l), thus indicating that none of the mercaptan and sulfides failed to be absorbed by the mercuric cyanide and benzene. Melting point determinations of the mercuric derivatives confirmed the identity of the methyl mercaptan, dimethyl disulfide, and dimethyl sulfide. LITERATURE CITED
(1) Ball, J. S.,U. S. Bur. Mines, Rept. Invest. 3591, 1 (1941). (2) Bell, R. T., and Agruss, hI. S., IND.ENG.CHEM.,ANAL.ED., 13, 297 (1941). (3) Borgstrom, P., and Reid, E. E., Ibid., 1, 186 (1939).
Challenger, F., and Charltori, P. T., J . Chem. Soc., 1947, 424. Claxton, G., and Hoffert, W. H., J . SOC.Chem. I n d . , 65, 333, 341 (1946). Faragher, If-. F., Morrell, J. C., and l l o n t oe, G. S.,I n d . Eng. Chem., 19, 1281 (1927). Kolthoff, I. M., and Sandell, E. B.. “Textbook of Quantitative Inorganic AnalyBis,” Kew YOIk. lIacmillan Co., 1948. Kolthoff, I. N., and Stenger, V. A , , “Volumetric Analysis,” Vol. 11, h-ew Tork, Interscience Publishers, 1947. Rlalisoff, W. AI., and Anding, C . E., Jr., ISD. ENG.CHEM., ANAL.ED.,7, 86 (1935). Sampey, J. R., Slagle, K. IT., and Reid, E. E., J . A m . Chem. Soc., 54, 3401 (1932). Seyfried, W. D., Chem. Eng. S e w s , 27, 2482 (1949). RECEIVED for review blarch 7, 1953 Accepted August 17, 1953. Journal Series Paper, N. J. Agricultural Experiment Station, Rutgers University, The Stat? University of Xew Jersey. Department of Microbiology. Investigation supported by a grant proi,icled by the Texas Gulf Sulphur Co.
Determination of Traces of Hydrogen Fluoride in Inert Gases D. L. RlANNING AND J. C. WHITE Analytical Chemistry Division, Y - I 2 Plant, Oak Ridge ,Yational Laboratory, Oak Ridge, Tenn. A method has been developed for the determination of traces of hydrogen fluoride in inert gas sweepings. The gases are passed through a dilute solution of boric acid and the increase in conductivity of the solution, which is produced by the reaction between boric acid and hydrogen fluoride to form a complex strong acid, is measured. The relationship between conductivity and concentration of hydrogen fluoride was linear for the range 0 to 1 mg. of hydrogen flnoride. As little as 10 micrograms of hydrogen fluoride can be determined by this procedure. The method is particularly valuable for industrial application when no other ionizable gases are present.
solution is :I strong acid (of the aanie order as hydrochloric acid), infinit,ely stronger than either boric acid of hydrofluoric acid (K1 a t 25’ C.) ( 2 ) . Hence, a solution of boric acid = 6.89 X will show a distinct increase in conductivity when hydrogen fluoride is added. -4study was made to adapt this reaction to the det,ermination of t,races of hydrogen fluoride in inert gas sweepings. EXPERIMENTAL
T
IIE detection and determination of traces of hydrogen fluoride in inert gases have been of recent interest’. The ideal method for this type of determinat,ion is one that could be conducted a t the site of the gas outlet, be relatively simple to accomplish, provide a rapid and accurate analysis, and be sufficiently sensitive to determine the gas in concentrations of the order of a few parts per million. Such limitations tend to Pliminate the known colorimetric methods for fluoride. -4 proposed method which would meet these criteria is a conductometric measurement of a scrubbing solution through which the gases have been passed. Possible solutions which would serve as scrubbers include solutions m-hich, contain cations that form slightly soluble fluoride salts such as calcium, barium, and lead, or solutions which contain an ion that forms complexes wit,h the fluoride ion. Water itself, because of t,he great solu1)ility of hydrogen fluoride in water, could be used as a scrubber. Boric acid, H,B03, was considered a promising scrubhing solution in this instance. A solution of boric acid has several inherent advantages with respect to ultimate conductometric determination; it is slightly ionized in water ( K = 6.4 X at 25’ C.) (1) and would hence have a specific conductance of the same order as water, and it reacts rapidly with fluoride ion to form a soluble complex acid. The reaction can be written: &Boa Wamser
+ 3HF +HBF30H + 2H20
(1)
(4) reports that monohydroxyfluoboric acid in aqueous
A standard, dilute solution of hydrofluoric acid x a s used to provide hydrogen fluoride. A stock solution was prepared by diluting 1 ml. of 48% hydrofluoric acid to 1000 ml. and stored in a polyet,hylene bottle. This solution was standardized by a conductometric titration against standard alkali solution. Known amounts of hydrogen fluoride were taken from this stock solution by appropriate dilution. The scrubbing solutions were prepared from a stock solution of boric acid which contained 10 grams of boric acid, C.P. reagent grade, dissolved in 1 liter of freshly boiled, distilled water. This solution v-as not standardized, but its concentration calculated to give 0.016-$1 boric acid. The experiments were conducted by adding from a buret knoFn amounts of hydrofluoric acid to 100 ml. of boric acid and nieasuring the conductivity of the solution after each addition. The final volume of solution in all cases was 110 ml. Conductivity measurements were made with a Leeds and Northrup conductivity bridge using a “dip-type” cell with a cell constant of 0.1 cm-1. Five concentrations of ecrubbing solutions were used: 0.16X, 0.10M, 0.01M, 0.0016M, and 0.0003M. No precautions were taken to regulate the temperature. The temperature control of the conductivity bridge was set a t the temperature of the dip cell. Inert gases, which were suspected to contain hydrogen fluoride, were sampled by passing the gases through the scrubbing solution of boric acid and then through a wet-test meter to record its volume. For continuous monitoring of the gas stream, the conductivity cell was located directly in the gas outlet line; for periodic monitoring, the scrubbing solution was removed and its conductance measured. Borosilicate glass containers were used and showed no evidence of etching. RESULTS A X D DISCUSSION
A tabulation of the results is given in Table I, which shows that the conductivity of the boric acid scrubhing solution increases n.ith the addition of hydrofluoric acid. As expected, the conductivity decreases with decreasing concentration of boric acid. .4larger increase in conduct,ivity is produced by the addition of 1007 of hydrofluoric acid to a solution of boric acid t,han
V O L U M E 25, NO. 11, N O V E M B E R 1 9 5 3
1649
for 1007 added to water alone. It niay be assumed that this increase is due to the formation of a stronger acid than either hydrofluoric or boric, thus producing the highly mobile hydrogen ion and increasing the conductivity of the solution. A comparison of the increases in specific conductance of the scrubbing solutions produced by the addition of hydrofluoric acid with the concentration of boric acid in the solution is presented in Table 11. The comparison indicates that a concentration range of 0.01M to 0.001.11 boric acid is optimum in order to obtain the maximum increase in conductivity per 1007 of hydrogen fluoride. No significant maximum between thrse concentrations was noted in subsequent experiments
18 00
c
R
0 161?
The data from Table I have been plotted in Figure 1. This plot of conductance against micrograms of hydrogen fluoride added shows graphically the relationships mentioned previously. The slope is essentially a straight line for the optimum boric acid concentration range noted. The recommended procedure is to establish a standard curve for a determined boric acid concentration by adding hydrofluoric acid and measuring the conductivity. These measurements plotted as shown in Figure 1 serve as a standard curve. The slope is also a measure of the order of sensitivity of the method. With a standard volume of 100 ml. of scrubbing solution, the concentration of hydrogen fluoride in the solution is approximately 1 X lo-* M following the addition of 1007 of hydrogen fluoride. The limit of sensitivity compares favorably with the most sensitive known colorimetric methods for fluorine. In actual practice it is more generally advantageous to express t.he value in terms of hydrogen fluoride per million parts of sweep gas. The use of a fairly sensitive conductivity bridge permit.s the measurement of as little as 1 X 10” M acid or 0.17 per ml. (part per million). The reproducibility of the measurement of the specific conductivity of 0.01 M boric acid was of the order of 10% between lots. For solutions containing enough hydrogen fluoride to produce a specific conductance of the order of 1.0 x 10” mho, the precision of the measurements was approximately 8%. The error of the method is k 1 0 y for the range tested, 20 to 1000y.
Figure 1. Effect of Micrograms of Hydrogen Fluoride on Conductivity of 100 MI. of Boric Acid Solution
Table I. Specific Conductivity of Solutions of Boric Acid Containing Traces of Hydrofluoric Acid Hydrogen Fluoride, y 0 10
20 30 40 50 60 7n ._ 80 90 100
0.16
M 8.86 9.50 10.3 11.1 12.0 13.1 14.0 15.1 16.5 17 5 18 8
Mhos X 107, Volume, 100 hI1. 0.0003 0 0016 0.10 0 01 M M M M 2.45 1.75 4.32 2.60 3.09 2.54 5.35 3.36 4.05 3.60 6.25 4 55 5.27 6 .OO 4.85 7.33 6.63 6.25 7.50 8.50 8.05 7.65 9.00 9.65 10.4 11.0 9.05 9.45 11.8 10.6 11.0 12.5 13.3 12.1 14.0 12.5 15 0 13.6 15.0 14.0 15.0 16.0 15.5 16.5
0 0 (Water) 1.02 1.55 2.40 3.46 4.45 5.55 6.65 7.85 8.94 10.1 11.3
Table 11. Increase in Specific Conductance of Solutions of Boric Acid by Addition of lOOy of Hydrofluoric Acid Solution, Molarity n 6.0003 0.0016
0.003 0.016 0.10 0.16
Conductivity, Mhos X 107 10.3 13.2 13.0 13.3 13.4 12.2
10.0
I
05 HYDROGEN
I FLUOR11
mg.
Figure 2. Effect of Milligrams of Hydrogen Fluoride on Conductivity of 100 M1. of 0.01M Boric Acid Solution
Further experiments were conducted using 0.01M boric acid to determine the extent of linearity of the slope with increasing hydrogen fluoride concentration (Figure 2). The relationship appears to be linear up to 5OOy of hydrogen fluoride. Beyond this concentration, the increase in conductivity per standard addition of hydrogen fluoride becomes smaller with increasing concentration of hydrogen fluoride. With a total of lOOOy of hydrogen fluoride present, the increase in conductivity per 1007 of fluoride is 5.0 X lo-’ mho, a factor of about 2 less sensitive than a t the 1007 level. The obvious limitation to the method is that the only ionizable species which are formed in the boric acid solution occur as a result of a reaction with hydrogen fluoride. This is not a serious
ANALYTICAL CHEMISTRY
1650 limitation for this particular application. The possibility exists that the conductometric measurement can be extended to the determination of traces of fluoride which can be liberated as aqueous hydrofluoric acid, as in the familiar pyrohydrolysis (5) technique. The actual ionic species in the boric acid-hydrogen fluoride solutions were not identified. However, Wamser (3, 4)has studied the system boron fluoride-water and has established that the rate of Reaction 1 is extremely rapid. He has shown further that hydrolysis of monohydroxyfluoroboric acid proceeds rapidly as HBF3OH and
+ H20Rapid +HBFs(0H)S + H F
+ H F +HBFi + H?O
the conductance of which was reproducible for only about 1 minute before a slow drift to higher readings became evident. The effect of the hydrolysis of a very dilute concentration of monohydroxyfluoroboric acid on its conductance is apparently insignificant, in which case readings are reproducible. As the concentration increases, however, this effect becomes increasingly pronounced and erratic readings are obtained. LITERATURE CITED
(2)
The reaction between hydrofluoric acid and monohydroxyfluoboric acid is extremely slow, however. HBFaOH
ductances of solutions of greater concentrations were not nearly
as stable. The extreme is illustrated by the conductance of a solution of boric acid, 0.5 with respect to hydrogen fluoride,
(4)
The results of this investigation substantiate the work of Wamser. The increase in conductivity of the boric acid solution upon the addition of hydrogen fluoride was noted immediately. The readings were reproducible for a t least 30 minutes for hydrogen fluoride concentration less than 1 X IO-' M , but con-
(1) Hodgman, C. D., "Handbook of Chemistry and Physics," Vol. 3-1, p. 1561, Cleveland,Ohio, Chemical Rubber Publishing Co., 1952. (2) Simons, J. H., "Fluorine Chemistry," T'ol. 1 , p. 252, Kevi York, Academic Press, 1950. (3) Wamser, C. A., J . Am. Chem. SUC.7 0 , 1209-13 (1948). (4) Ibid., 73, 409-16 (1951). ( 5 ) Willard, H. H., and Winter, 0 . B., ISD.ENG.CHEM.,ANAL.ED., 5, 7 (1933). RECEIVEDApril 30, 1953. Accepted August 10, 1953. The Oak Ridge National Laboratory is operated by the Carbide and Carbon Chemicals Co.. a Division of Union Carbide and Carbon Corp., for the Atomic Energy Commission. The work reported here was carried out under Contract No. W 7402-eng-26.
Curve for Argentometric Determination of Cyanide JOHN E . RICCI Department of Chemistry, Y e w York University, New York, iV. Y . The shape of the curve of p[Ag+]in the titration of potassium cyanide with silver nitrate has been investigated, to determine the position of the inflection point relative to the equivalent point. The relations considered include the precipitation of silver cyanide and the effect of ammonia upon the titration. Equations are derived for the calculation of the equilibrium cmnstants from the characteristics of the titration curve.
I
N AN earlier article ( 8 ) on the interrelations of the equilibrium conetants in aqueous solutions saturated with silver cyanide, the relation of the point of precipitation of silver cyanide or of silver iodide, as titration end point, to the equivalent point in the argentometric titration of cyanide was considered. The present article is concerned with the shape of the titration curve. or the curve of p[Ag+] against the quantity of silver nitrate added, in order to determine the position of the inflection point in the curve and its relation both to the equivalent point and to the pomt of first appearance of a precipitate. Two sets of conditions are considered: the titration of pure aqueous potassium cyanide with silver nitrate, in which the hydrogen ion concentration is variable during the titration; and the titration of potassium cyanide in presence of excess of ammonia, during which the value of [H+] is practically constant, fixed mainly by the ammonia. -4few of the expressions involved are identical with those of the previous discussion; two errors that crept into that paper as published are corrected here. For ease of comparison, all symbols remain as before, with the exception of the use, here, of A in place of K. for the ionization constant of hydrocyanic acid. NOMENCLATURE
el
= analytical concentration of potassium cyanide = analytical concentration of silver nitrate
z
= analytical concentration of potassium iodide = [A@;+]
c
b e,
= analytical concentration of ammonia
Y 2
H W A Kb
[H+][OH-] = [ H + l [ C N - ] / [ H C N ]= 4.0 X [NH,+][OH-]/[NHa] 1.8 X [Ag+][CN-]'/[Ag(CN)2-] = 1.4 X [Ag+][NH3]'/[Ag(NHa)2+]= 6.0 X IO-* [Ag+][CN-] a t saturation with AgCN, = 1.2 X lo-'* xy a t saturation with AgCN, = P 2 / K = 1.03 X lo-'* [.4g+]/I-l a t saturation with A4gI. = 8.5 X
K K' P P' p3
n m pz # 4'
= =
r s
= 1 H/A = P' 4r P =
4 u
=
Q
y g
-drnjdpz
+
= -
v
( ( (
=
= =
)b
)p
= = value a t equivalent point = value a t inflection point
value required for precipitation activity coefficient of a univalent ion = ionic strength
= =
The italic capitals represent mass constants, and activity constants are in bold face. The e uality sign ( = ) is used only for relations which are exact by dexnition, and the symbol Z is used for all approximations, even if very accurate. c and cl represent not initial concentrations or concentrations
