July, 1961
NOTES
1279
to stand for 1,wo hours a t room temperature. I t then was neutralized to the brom cresol green end-point with perchloric acid, and a solution of lead nitrate was added. No lead sulfide was formed, thus indicating that the formation of sulfide ions by the alkaline decomposition of thiourea was negligible under thew condit,ions. Analyses.--Hydroxide, lead6 and thiourea' were d e k mined by standard procedures. The hydroxide ion concentration of the samples was determined by back-titrating the excess perchlorio acid quenching solution with sodium hydroxide to the brom cresol green end-point. Lead was titrated (in the presence of tartrate ions and ammonium hydroxide) m-ith disodium ethylenediaminetetraacetate to the eriochrome black T end-point. The thiourea was treated with bromine (generated by sodium bromate, sodium bromide and hydrochloric acid), t,he excess bromine was removed with potassium iodide, and the iodine thus formed w m determin'd by titration with sodium thiosulfate.
obtain the reaction orders from the rat,es at corresponding stages of the reaction and from the half times of the reaction. The orders with respect to lead, thiourea and the hydroxide ion were found to change as the reaction progressed. In addition, solutions prepared from different batches of sodium hydroxide reacted a t anomalously different rates, and the reaction rates of the solutions decreased with the age of the sodium hydroxide solution from which they were prepared when t,he latter had been stored in Pyrex, but not when it had been stored in polyethylene (see Series 7). In the second group of runs of Series 7 , it was found that stirring decreased the average reaction rat'e: in the solutions stirred for 0.3 min., the reResults and Discussion action was 51% complete after 14.3 min., while From the a,na,lytical determinations of Series in those stirred for 12 min., it was only 32-3370 1-3, the ratio of The iiuniber of moles of hydroxide complete. The samples taken from the lattaerwere ion to the number of moles of lead consumed during markedly more transparent, indicating that conthe reaction n7as calculated. I t s value was con- siderably less colloidal lead sulfide was present,. sistently a little below 2.0, but by an amount that These results can be explained by the hypothesis (for most runs) was within the experimental error. that the lead sulfide formed by the reaction is The value was independent of the extent to which act'ing as a surface catalyst and that stirring assists the reaction had .proceeded. From the analyses of in its coagulation, thus reducing its effect,ix-esurface Series 4, the ratio of the number of moles of thiourea area a,nd therefore its catalytic effect. to the number of nioles of lead consumed was calAcknowledgment.-The aut'hor wishes to exculat,ed to be a little below 1.00 for both runs. The press t,hanks to Dr. Robert F. Brebrick for his help. value was within the limits of the experimental ful advice and encouragement. error for the run having the lom-er initial lead concentration and just out'side theFe limits for the TBIS IONIZATION CONSTANT O F other run. The values were independent of the HYDROXYLSMIKE extent to which the reaction had proceeded. These ratios help to confirm t~hefollowing equation which BY R. A. ROBINSON AND T'. E. BOWER had been postulated previously by several au.Vational Bureau o f Standards, WashingLon, U . C . thors3-5 tleceirid February 17, 1061 I'b-r+ + (X€I,)&S + 2 0 H - = PbSJ + HcNCN + 2HzO Although hydroxylamine has heen known for Bruckniann, ;Sahasrabudhey and &all4 and nearly a century, its ionization constant is not) Pick,2 suggested that the reactioii involved t,he well determined. The conduct,ance work of Ross' formation and brcnkdown of a lead-t>hiourea inter- gives pK 6.1 at 18' using his data a t higher coiicenmediate compound, while Whitcher3 believed that traLions (-0.2 IIr) ; the colorimetric measurements it involved the formation of sulfide ioiis by the of Olander* at-20" give pK 6.09 while there are t'hree alkaline decomposition of thiourea. The negligible detially 1ieutr:tlized by :tddit,ioii of sodium hyruns of Serits 5 completely removed t'he induction droxide so that a port,ion is present as t,he hydroperiod ( L e . , thc initial period of slow rate). In chloride, S, and a port,ion as the frcle base, R, and if the average rate during t,he solution contains also a "Colored" indicator, Series 6, it \vas found tmliatm the first 5.2 niiu. n-as roughly proportional to the such a,s ~~,~-dinitrophciit~I, the equilihriii are given weight of thc 1c:td siilfidc powder added. E'rom hS t,hese observations it, was concluded that. the colp K = p H - log (11 - D i i i ( L ) y - 1)) - log -1 loidal lead su1fid.e formcd hy t,he reaction was for the indiwtor aiid uctirig as a eurfa,cc cat,aI . As the amvuiit of pKa = pII - l~gjB]/jS] + log - j lend sulfide grew, t hc reaat~ionrate increased until __.___ ~the lowered concentrations of the reactants and/or ( I ) W. H. Ross, Proc. Trans. Nova Scotian Insl. S a . , 11, 95 (1905). the coagulation of the catalyst reversed the trend. (2) A. Olander, 2. physik. Chem., 129, 1 (1927). (3) K. Winkelbleoh, ibid., 36, 546 (1901). Using t,he "lead concent'ratjion us. time" curves F. Ishikawa and I. Aoki, B7dL Soc., I n s f . Phys. Chem. Res. Tukyo, of the runs of Series 1--3.an athempt was made to 1 9 (4) , 136 (1940). 6 ) .l. 1,. Piiih.iiiii : i n , c'. 1'. lieliner. . I n n i . C k e m . . 27, 446 (195,;). 17) .iinc*rican Cy:tnainiil Coiiir)any. 'leclinicd Hats Slreet '1'0110,33-02, Bound Brook, N. J. ( n o dat.e).
( 5 ) 11, Iiagisawa. ibid., 20, 2.51 (1941). (ti) "Stability ~ ~ u l i s l i l i i t s .I'nrt " 11, 'rtic Clleriiicnl 10.58, [I. 6 1 .
sc
NOTES
Vol. 65
TABLE In IONIZATION CONSTANTS OF HYDROXYLANINE AT 20, 25 A N D 30" ISl?
,;.
[3$
1B1,
mole/l.
mole/l.
Indicator, mole/l.
D1
c
200
*
25'
7
30'
200
2%
30'
0.2834 0.1452 1.952 6 . 8 x 10-6DNP 0.027 0.024 ... 0.924 0.925 ... 0.3238 0.0732 4.423 5.6 X DNP 0.025 0.023 ... 0.764 0.768 ... c 0.3430 0.1095 3.132 5 . 6 X 10-6DNP ... 0.022 0.022 ... 0.761 0.761 D 0.1061 0.1074 0.988 1 . 7 x 10-4NCP ... 0.090 ... ... 0.717 ... S = hydroxylamine hydrochloride, B = hydroxylamine, D N P = 3,4-dinitrophenol, NCP = 2-nitro-4chlorophenol. The four stock solutions contained S and B a t the concentrations recorded and sodium chloride a t the same concentration as B; z ml. of stock were made up to 100 ml. The diluted stock contained indicator a t the concentration recorded. T o obtain D1 and Dz,solutions of the same indicator concentration were made in 0.01 N HCI and 0.01 N SaOH, respectively. The following values of D were obtained:
A=! B'f
Series
0
A
C.
B C
D
10
15
20
25
50
0.656
0.682 0.671
0.401 0.458 0.431
0.420 0.471 0.449
0.688 0.678 0.444 0.430 0.481 0.456
0.699 0.688 0.448 0.438 0.491 0.462 0.277
0.708 0.694 0.456 0.442 0.498 0.469 0.283
0.721 0.711 0.479 0.461 0.512 0.485 0.297
x 4 5
...
20 25 20 25 25 30 25
... ...
...
...
...
Measurements in series A, B and C wcre made at 400 mp and in series D a t 420 mp. The average pR, values wcre: 25" 30 Series 20 O 5.96 ... A 6.03 5.97 ... B 6.05 C ... 5.95 5.84 n ... 5.93 ... O
where K a is the acid ionization constant of the hydroxylammonium ion. D is the optical density of the solution, D1that of a solution of the same stoichiometric Concentration of the phenol containing sufficient hydrochloric acid to give pH -2 and Dzthat in a solution containing sodium hydroxide a t pH -12; y is calculated as -log y = Ad?/(l dI)- 0.21, Hence P K A = pK - log lBI/[Sl log (D - Di)/(Dz - D )
+
+
+
2 log
Y
The first term on the right already has been measured' a t 25"; it has been redetermined and the value of pK 5.42 confirmed. The second term is known from the composition of the solution and the third is found by spectrophotometric measurements. Experimental Hydroxylamine hydrochloride (Fisher's Certified Reagent) was recrystallized from aqueous ethanol and dried in uacuo over "Drierite." The drying was a very slow process, not complete even after six weeks when electrometric titration gave an assay of 99.8%. Measurements were made with a Beckman Model D U instrument, the cell compartment being maintained a t constant temperature, rtO.O5O, by circulating water froin a thermostat. Three series of measurements with dif'ferent [B]/[S] ratios were made at 25' using 3,4-dinitrophenol as indicator and a fourth series was run with 2-nitro-4-chlorophenol as indicator. The p K of the latter has been measured recently.* Details of the meaeurements are given in Table I. Much difficulty was experienced from the instability of the solutions and no reliable result could be obtained from a solution which was more than a few hours old; the instability was more marked with solutions of low concentration and high pH. The ionization constants of 3,4-dinitrophenol a t 20 and 30' were needed; measurements at these temperatures are reported in Table 11. With these data, further spectro(7) C. M. Judson and M. Kilpatrick. J . A m . Chem. Soc., 71, 3110 (1949); R. A. Robinson. M. M. Davis, M. Paabo and V. E. Bower, J . Reaearch N o € . Bur. Standards, 64A,347 (1960). (8) V. E. Bower and R. A. Robinson. J . Phys. Chem., 64, 1078 (1960).
photometric measurements (Table I) gave the ionization constants of hydroxylamine a t 20 and 30".
The ionization constants of 3,4-dinitrophenol are: 20°, pK 5.46, 2 5 O , pK 5.42; 30°, pK 5.38. The average values for hydroxylamine are: Z O O , p K a 6 . 0 4 ; 25O, P K a 5.96; 30°, p K a 5.84. IONIZATION
CONSTANTS
Buffer
TABLE 11' 3,4DINITHOPHENOL AT 20 30"
OF
D
AND
Ph-
20°, 400 mp, D1 = 0.025, D, = 0.764 f 0 394 5.47 fi ,416 5 46 11 440 5.44 30°, 400 mp, D1 = 0.022, Dz = 0.761 f 0.427 5.38 R .440 5.39 h ,467 5.37 a The molarity of 3,4-dinitrophenol was 5.6 X throughout. Buffer f, g and h were mixtures of x M sodium hydrogen succinate and XM sodium succinate where 5 = 0.05, 0.025 and 0.01, pH 5.348, 5.406 and 5.477 at 20°, pH 5.343, 6.403 and 5.474 at 30" and -log y = 0.117, 0.102 and 0.077 a t 20" and 0.119, 0.104 and 0.078 at 30°, respectively.
A comparison of its thermodynamic properties with those of the ammonium iong gives the following data: Hsdr?xylammonium
Ammonium ion
ion
AGO 34,000 AHa 34,000 AS" -6 to +8
52,780 j. mole-' 52,200 j. mole-' 1.87 j. deg.-l mole-'
-
The basic ionization constants of hydroxylamine are: 20°, p K b 8.13; 2 5 O , p K b 8.04; 30°, pKb 7.99; so that hydroxylamine is a very much weaker base than ammonia. (9) R. G. Bates and G. D. Pinching. J. Research Natl. Bur. Standards, 42,419 (1949); J. Am. Chem. SOC.73, 1393 (1950).
