THESYSTEM SOD~UM FLUORIDE-HYDROGEN FLUORIDE-WATER
April 20, 1959
free C1- ions in our system. If we assume t h a t the LiCl exists in solution in an undissociated form, then by the reaction LiCl BuOLi BuOLi,+ C1- more chloride ions are available for exchange. Still there is no evidence available for the existence of complexes of the type t-BuOLiz+. We conclude t h a t i t is hard to offer an explanation for the t-BuO- effect. The eflect of acetate ions on the rate of chlorine exchange is not clear. It is hard to assume acetyl hypochlorite as intermediate because its rate of formatioil is slower8 than the catalytic effect of acetate ions on the chlorine exchange. The pH dependence here is not well established, whereas the dependence on the chloridc ion conccntration points tc competing reactions between AcO- and C1- ions. Wc may conclude that the effect of acetate ions on the chlorine exchange requires further investigation.
lytic effect of small amounts of water on the rate of exchange, points to an effect of ion pair dissociation. I n the acid region we pick up a reaction path involviiig acid catalysis, but we see t h a t the presence of an acid is no prerequisite for the KOC1C1- chlorine exchange, as i t is in the case of the HOC1-C1- exchange. T h e acid path becomes appreciable only a t acid concentrations above molar where we may assume t-BuOCl+ as interinediate. Estimating the dissociation constant of t-BuOHCl+ a t 10' we derive a specific rate constant for this reaction k = 4x1071. mole-' min.-l a t 27". I n the basic region we are surprised by the well established catalytic effect of t-BuO- concentration on the rate of exchange. We suggest three ways to explain this effect: one is to assume t h a t a butoxide ion forms a complex with t-BuOC1 which facilitates the interaction with C1-. I t is hard however to see, why should (t-BuO)rC1- be more liablc to a nucleophilic attack than t-BuOC1; if such a complex would be formed, i t might probably suppress the rate of exchange. An alternative explanation may be b y the reaction f-BuOC1 t-BuOLi t-BuOCl+ tBuO- whereafter
+
Li
+
t-BuOCl+ - . is attacked by C1-; Ll
t-BuOCl+ _.
+
Ll
1821
+
+
Conclusion Our results give mole evidence that the HOCl system difiers basically from the HOI3r system. Having a much stronger HO-X bond, HOCl follows different mechanisms of interaction with nucleophilic reagents the HOCl molecule requires occasioiially an additional proton to make its chlorine accessible. Next it has been shown t h a t the trihalide like ions HOX-2 or HOXYmay occasionally occur as intermediates in HOX reactions.
+ C1-
t-BuOLi C1,. This assumption would require a second-order effect of LiCl concentration, which has not been observed. The third explanation may be the effect of BuO- on the availability of
REHOVOTH, ISRAEL -
[CONTRIBUTION FROM T H E CHEMISrRP
DEPARTMENT O F T H E AGRICULTURAL AND
The System NaF-HF-H,O
at 0 and
BY JAMES S. MORRISON AND ALBERT W.
M E C I I A N I C A L COLLEGE O F T E X A S ]
- lSol JACIIE
RECEIVED JULY 7, 1958 The system NaF-HF-H20 has been studied at 0 and -15". The solid phases include S a F . € I F , SaF.2HF, NaF.3IIP' and NaF.4HF at both temperatures. N a F exists in the 0' case and presumably at - I s o .
cerning the fluorides in this system. Siiiions7 Introduction Ternary systenis involving an inorganic fluoride, states that, with aqueous or anhydrous HF, the hydrogen fluoride and water have been studied by acid salt Tu'aF.HF results and is the only acid relatively few workers. Tananaev has investi- fluoride of sodium known. According to Jache and gated systems involving hydrogen fluoride, water Cady,s the analysis of the solid phase in equilibrium and lithium fluoride,2 potassium fluoride, zir- with H F saturated with N a F indicated the presence conium(1V) fluoride4 and iron(II1) f l ~ o r i d e . ~of 1 mole of NaF to 4.10 of H F a t -24.3". TanaClarks has listed some data on the solubility of naevg has also studied part of this system but his sodium fluoride in various concentrations of results seem not to have been noted by later workers. aqueous hydrogen fluoride. R'e have studied this system a t 0 and -15" Our interest in the ternary system sodium fluoride-hydrogen fluoride-water (NaF-HF-H20) was using the Schreinemakers wet residue method.1° stimulated by the lack of complete data for this Experimental system and the disagreement in the literature con(1) Presented a t t h e 133rd National Meeting of t h e American Chemical Society, S a n Francisco, California. (2) I. V. T a n a n a e v , Kkirri. K e d k i k h . Elctrit.,ilui,. d k a d . S n i i k . S.S. S . K . Z i i s l . 0 b s h t . i . .\-coI-~. K l i i v i . , 1, 33 (1954). (3) I. V. T a n a n a e v . J. APPI. Chew. ( U . S . S . K . ) , 11, 214 (IW49). (4) I. V. T a n a n a e v . N. S. Nikolaev a n d Y. A. Burlaev, Z / f r t r , Seorg. K h i m . , 1, 274 (1956). ( 5 ) I. V. T a n a n a e v . J . A p p l . Chem. ( U . S . S . R . ) ,19, 1018 (1946). (6) George L. C l a r k , THISJ O U R N A L , 41, 1477 (1919).
Samples were made up from reagent grade N a F arid aqueous H F . IYhen HF of concentration greater than 48% was needed it was made up from mixtures of anhydrous IIF
(7) J . H. Simipns, "Fluorine Chemistr)-," 1.01. I , Academic Press, I n c , S e w I'ork, N. T., 19.50, p. 27. (8) A. W. Jache and G. 11, C a d y , J . P h y s . (,.iie?,r,, 66, 1101'~(19>2) (9) I . V T a n a n a e v , J . Geir. Chrrii. (C..S'.S.K.),11, 2fj7 (1941). ( I O ) F. E. Purdon a n d 1:. W. Slater. "Aqueous Solution and t h e Phase Diagram," Arnold and C o . ,London, l!HG.
S.ATORRISON
1 S??
JAZ~IICS
AND
ALBERTW.
VOl. 81
JACHE
TABLE I1
7
IVct re5idue.;. - - w t . I,;--€11: "at'
-
ANALSSIS OF Co.nposition of solld phase
-wt.
TIlE SOLUTIONS AND 1b'ET
Solutions
%--
Wet residues. 7 - w t .
%-
NaF
HF
...
...
10.78 13.03 17.11 24.09 23.22 25.57 28.59 29.36 34.09 36.34 37.03 1n.10 41.83 41.03 44.29 41.75 49.46 48.22
28.97 32.91 23.21 39.03 29.24
50 25 tj2.87 54 23 54.86
3.80 2.02 1.98 0 . 603 .5S7 ,567 ,650 ,754 1.09 1.68 2.37 2 47 4.59 5:30 5 18 8.37 11.37 I*'] . 11 1,574
55 68
1;i 75
5ii.1:1
13 90 17.30 17.34
HF
0 0.365 2.82 11.18
13.71 17.04 20. ox 24.88 29.15 :14.48 38.15 38 95 15.03 46.10
4646
SS 58 59 27 59 97 lil .49 (il 915
lW.77 69.2(? 71.1.3 75.10 75.23 78.21 81.02
17.13 17.01 17 08 14.16 12.29 12.70 15.12 14.35 17.32 19.018
$52'.Ti 53. 88 54.95 59.G9 63.02 59.74 li3.20 (is.57 87. 13
NaF
RESIDUESAT -15" Composition of solid phase
31.81 30.40 24.43 22. '77 21.04 29.67 28.85
21.77 27.53 29.55 40.54 24.45 34.41
07.00 89.9li 70.38 72.32
33.89 26. 4li 30.08 30.73 27.65 30.35 26.26 24. 6.3 25.53 26 47 27.45 24.77 26.01
..
..
regulated t o k O . O G o . These bottles were agitated in the b:tth. If n o solid phase remained, mnre NaF was added. Liquid sarnplcs wcre removed by means of a sampler using = li: liypixlerniic ncedlc fitted through a Teflon adapter t o siii;+ll pJvlyc.tliyleiie bottle.
i
\ \\ \ \
\\\\\\
YYM
Fig. 1.-The system NaF-HF-HQO a t 0'. L'xrL~iis niisturcs o f IIF, IIzO a i i d Nal: were p1:icetl i n 250 cc., wide r i i o u t l i polycth;.lene Ixttles. These were p l x e d in ail ctliylcne glycol-water bath whose temperature was
Fig. 2.-Thc
system NaF-HF-H20 a t -13'.
The liquid samples and wet residues were analyzed for HE' by a simple titmtion with sodium hydroxide solution.
T h e NaF was determined by careful evaporation in pl'ttinum crucibles over a heat lamp followcd by ignition at 750'. T h e weighing form was NaF.
Discussion and Results The results of the analyses of the solution and wet residues are given in Tables I and I1 and have been plotted in Fig. 1 and 2. At - 15" a t H F concentraall phases appear to be solids. tions of less than ll~o, The solid phases existing in equilibrium with solution a t both temperatures include N a F . H F , NaF.2HF, N a F . 3 H F and NaF.4HF. N a F exists
[CONTRIBUTION FROM TIIE RESEARCIi
in the 0" case and presumably a t - 13". N o solid phases containing water exist. The H F (wt. yo) concentration limits for the solid phases a t 0" are: NaF, 0 to 9.5c7,; N a F . H F , 10 to 577,; NaF.2HF, 5 3 to 5S.5yo; NaF.3HF, 3.5 to 6256; and NaF.4HF, G2 to 100yo. In general, the solid phases containing H F are more stable a t lower H F concentrations at - 15" than they are a t 0". The solubility of N a F is in general sonicwhat lower at the lower temperature. COLLEGE STATION, TEXAS
LABORATORP, OLIN
hfATIIIESON CHEMICAL CORPORATION]
Change in the Ratio of Hydroxyl Groups Attached to Silicon and Aluminum Atoms in Silica-Alumina Catalysts upon Activation BY H. G. WEISS,J. A. KNIGHT AND I. SIIAPIRO RECEIVED OCTOBER18, 1958 A method for measuring the number of Iiydri)xyl groups attached t o silienn relative t o those attached t o aluminum atoms in silica-alumina catalysts is described. Essentially, the method consists of two steps: ( a ) exposing dibrirane to the catalyst in order to replace the pri)ton of tlic 1iydri)xyl group with boron atcms and ( b ) nica~uringthe nuniber o f these boron atoms t h a t can participate in bornn exchange with labc.lcd diborane. By this technique, it is found that initially hydroxyl groups are attached to both silic~~ii and aluiiiinnm a t i m s ; however, upon aging of tlic catalyst a t temperatures below 400" hydroxyl groups, lost as water, must conie predomiiiantly froin the silicon ratlicr than thc alurninuin atoms. At ternperatures above ca. 400" t h e ratio of hydroxyl groups attached to alumiiiuiii and silicon atoms goes through a rnaximuni value.
T h e relation of bound water to the surface structure of silica gel, silica-alumina and aluminuiii oxide catalysts has been the subject of several investigations.'-$ The influence of sinall amounts of water on the catalytic activity of alumina and silicaalumina also has been reported."'j In this connection various n i e t h o d ~ have ~ , ~ beeii devised for measuring bound water content of catalysts; however, such methods (ignition, D 2 0 or D 2 exchange, etc.) are satisfactory for bulk determinations but do not lend themselves readily to determination of the location of the bound water. The nature of bound water (hydroxyl groups) in silica gel has previously been elucidated by a studv of its reaction with ~ I i b o r a n e . ~Froin the relative :tinourit of dihorane consuinetl to hydrogen generated, it was tletermiried that the hydroxyl groups in silica gel lie principally on the surface of the solid, with the surface composition consisting of Si-0BHz units. Similar results were indicated fur reaction of diborane with silica-alumina. T h e diborane-treated silica-alumina catalyst'O exhibited a pronounced activity toward cyclization of acetylene to benzene, and from a detailed study of the cycliiation reactionll there were indications that alumina provided the active sites in silica(1) I. Shal'iro a n d I. I f ,F;-looo) there was a decrease in the hytlrogen-t1ihc)rane ratio, thus precluding the possibility of oiily ?.IOBII:! units forming. Boron-10 exchange experiments betweeii gaseous diborane anti the boronsurface coinpounds showed differences in the amount of exchangeable buron 011 the surface. Total boron exchange (1OOz) occurred in the borane-silica compound formed froiii diborane and silica gel. V-ith alumina, however, only one-half (50Ti8)of the attached boroii atoins werc capable o f cxchange. \\.hen exchange experiments were carried out in the silica-aluiiiina-borane coinpound, the amount o f boron participation varictl hetweeii these limits. From the equilibrium exchange values, the ratio of hydroxyl grouljs on aluniiiia to those on silica has been calculntcd. 'The results indicate that the hydroxyl groups in silica-alumina are mi both altiniinuni and silicon atoms and that those attached to silica are inore easily reiiiovetl on heating. Details of the exchange reactions between the diborane surface coniplexes and gaseous diborane are given.
