V O L U M E 2 7 , NO. 11, N O V E M B E R 1 9 5 5
1683
effective in either the presence or absence of oxygen. By using acetone as a solvent and suppressor (in the absence of oxygen), cuniene hydroperoxide can be determined satisfactorily in a colorimetric procedure based on the osidation of ferrous iron ( I d ) . Hon-ever, this procedure does not give accurate results n-ith peroxides in fats and oils. The addition of a suppressor can force the reaction between a peroxide and ferrous iron to proceed in the molar ratio of 1 to 2, as expressed in Equation D. I n mother approach, tlie addition of a monomer has been iised to force the reactioii lietween a peroxide and ferrous iron to proceed in the molar ratio of 1 to 1, a? expressed i n Equation G :
2ROOII
can serve for the determination of one of these reactants h y means of the other. LITERATURE CITED
Barnard, D., and Hargrave, K. R.,A n a l . Chim. Acta, 5, 476 (1951). Baaendale, J. I$., Evan.?, 31. G., and Park, G. S..T r a m . Faraday Soc., 42, 155 (1946). Benson, Clara, J . P h y s . C‘heni., 7, 1, 356 (1903); 8, 116 (1904). Bolland, J. L.. Sundralingam, -4.. Sutton, D. A , , and Tristram, G . R., T r a m . Inst. Rubber Ind., 17, 29 (1941). Bray, W.C., and Ilainiey, J. B., J . Am. Chem. Soc., 55, 2279 (1033). Fordham, ,J. W. L., and \Tilliams, H. L., J . Am. Ciiem. Soc., 73, 1634 (1951). and Keiss, J., Saturwissenschaften, 20, 946 (1932) ; Haber, F., Proc. Rou. S O C .( L o n d o n ) , A147, 332 (1934). Hartman, L., and White. 11. D. L., .ISIL.C H m f . , 24, 527 (1952). Kessler, F., Pogg. ~ u L . 119, , 218 (1863); 95, 224 (1855). Kharasch, 31. S . , Fono. I.,and Sudenberg, W., .J. Ora. Chem., 15, 703 (1950). Koithoff. I. 11.. and Carr, E. AI,, . ~ N I L . CHEM..25, 298 (1953). Kolthofi”, I. lI.,and Lairiiien, €I A , private communication, cited in (f4). Kolthoff,I. 31 , and 3Ieda!ia, A. I., J . Am. Cheni. Soc., 71, 3777, 3784. 3789 (1940). Kolthoff, 1. AI., and IIedalia, A. I.,ASAL. CHEM.,23, 59; (1951). Kolthoff. I. 31., and Stenger. 5’. A, “Volumetric Analysis,” vol. I, 2nd ed., p. 168, Interscience, Kew York, 1942. and Selson, J. S., ISD.ENG.C m h r . , h s a ~ . . E ~ . , Laitinen. H. -I., 18, 422 (1946). Lang, IZ., and Zaerina. J., 2 . anorg. ZL. a2lgem. C h ~ m . 170, , 389 (1928). Lea. C. €I.. J . Soe. Cham. Irid., 64, 106 (1945). Lee, T. S., in Friess, S.L., and Weissberger, .4., eds., “Technique of Organic Chemistry,” vol. VIII, pp. 100-30, Interscience, S e w York, 1953. Lenssen, E., and Lowenthal, J., J . prakt. Cheni., 87, 193 (1862). Livingston, Robert, in Friess, S . L., and Weissberger, A, eds., “Technique of Organic Chemistry,” vol. VIII, pp. 219, 224, Interscience, Kew York, 1953. Luther, It., and Rutter, T. F., 2.anorg. Chem., 54, 1 (1907). Luther, R.,and Schilow, S . , Z . physik. Chem., 46, 777 (1903). A n n . Chem. Justus Liebig, 325, 93 (1902). Manchot, W., Ilanchot, IT., and Wilhelms, O., Zbid., 325, 105 (1902). Nedalia, A. I., unpublished experiment.s. hlers, J. H., and Waters, W.A., J . Chem. Soc., 1949, Sl5. Schonbein, C. F., J . prakl. Chem., 75, 108 (1858). Schonbein, C. F., P o g g . Ann., 100, 34 (1857). Siggia, S.,ANAL.CHmr., 19, 827 (1947). Stein, G., and Weiss. J., J . Chem. Soc., 1951, 3265. Wagner, C., and Preiss, W., Z.anorg. u. allgem. Chem., 168, 265 (1928). Wagner, C. D., Smith, R. H., and Peters, E. D., ANAL.CHEM., 19, 976, 982 (1947). Westheimer, F., Chem. Rev., 45, 419 (1949). Wheeler, D. H., Oil & Soap, 9, 89 (1932).
+ 2 F e + + + 2 n ~ I I z = C “ X --+ 2 F e + - + 20II- + R O [ C H Z C H S ] ~ , O R ( G )
The meclimism of the over-all Iicwtiori G is of course the f a niiliur free ritdical chain l,ul!-niei,izatioii, n-ith initiation hy Step 18 follo\\-etl by:
RO.
+ CH?===CIIS
then propagntion: ROCHXHX ( n - l)CII,=CHS
+
ROCHzCHX
--+
(22)
RO(CH1CHX),.(23)
and termination: 2RO( CI1,CHX)n. +RO((’H,CHX)z,OR
(24)
Thus, the reaction betiveen cumene hydroperoxide and ferrous iron has been made t o proceed quantitatively according t o Equation G by the use of :icrylonitrile as the monomer ( 6 ) i n the absence of oxygen. From the standpoint of the acrylonitrile, Equation G i’eixeserits a n induced chain reaction, since the polymerization of acrylonitrile is a chain reaction ivhich is induced b?- the peroxideiron reaction. However. this chain reaction (Steps 23 and 2-1) does not involve either of the primary reactants (the peroxide and the ferrous iron’. Froin the st:rndpoirit of the prim:ii,~. reactants, the over-all induced Itencation G is simpl), a C O \ i l J i C Y i reaction, with a coupling index of unity. From the practical analytical standpoint, this illu.+;ttes a i l important difference between coupled reactions anti induced chain reactions. I n the absence of acrylonitrile, the over-all reaction consists of the primarj- Reaction D and the iiiduced chain Reaction E, which, because they proceed in an indefinite ratio, give an indefinite stoichiometry. I n the presence of Rufficient acrylonitrile, the primary Reaction D is completely displaced by the coupled Reaction G , which has a definite stoichiometry Jvith regard t o the primary reactants and therefore
RECEIVED for review July 28, 1955. Accepted August 2 6 , 1955.
8th Annual Sammer Symposium-Role of Reaction Rates
Kinetic Behavior of Halide Complexes H. M. NEUMANN Department of Chemistry, Northwestern University, Evanston,
The stepwise hydrolysis of halide complexes is discussed €rom the kinetic point of view covering the following subjects: experimental methods for studying the kinetics, structural factors affecting the rate, and mechanistic features of the reaction. Examples are also given of the utilization of kinetic data in identifying ionic species in solution.
111.
HEX metal halides are dissolved in the corresponding halogen acid it is a common reaction for the metal to become part of a complex halo anion. Isolation of salts containing the halo anion has been performed in a sufficient number of cases to leave little doubt as to the great prevalence of this type of reaction. Evidence t h a t such ions also exist in solution has been obtained by several different physical methods. It is only
1684
ANALYTICAL CHEMISTRY
the most electropositive elements that do not seem to form complexes of this n p e . It is also a well-known fact t h a t if such solutions are diluted the halo complexes generally decompose, and in many cases a precipitate forms. The nature of the precipitate depends on the metallic element; it may be a basic salt, hydroxide, or hydrous oxide, and often its actual nature is difficult to determine. I n accordance n i t h the modern conceptE of complex ions it is believed that the transition from the halo complex to the final product must proceed in several discrete steps, the initial steps of which are successive replacements of the halogen atoms in the complex. I n a few favorable cases the stepwise hydrolysis pioducts have been isolated. For example, in the case of chromium(III), salts containing the anions, CrC16--- and Cr(H20)C1,--, hare been isolated. With platinum(IV), salts containing the ions, P t C k - , Pt(OH)ClS--, Pt(OH)nCl,--, etc., are knoivn. These two examples illustrate that the hydrolytic species, whether present in solution or isolated as salts, d l be of a form where halide ions have been replaced by either a n-ater molecule or hydroxide ion. Which of these forms appears depends merely upon whether the ion M(H20)X,-” is a strong acid or weak acid under the conditions of the experiment. Presumably the case might also arise Rhere the acid strength could be such that comparable amounts of AI(H,O)X,-” and lI(OH)Xv-(2+1) n-ould be present under the experimental conditions. Although only a fem good examples of this stepwise hydrolvsis can be demonstrated by chemical isolation, it is assumed that this is a general type of behavior. I n some favorable cases, like that of the complexes of antimony(\-) in hydrochloric acid ( 7 ) , it is possible to demonstrate this type of behavior from physical el-idence even when isolation is not possible. From the kinetic point of view each step in the hydrolysis is merely a particular substitution reaction; hence some of the general features of substitution reactions are tvorth recalling. Taube (13), in reviewing these reactions, has divided complexes into two classes on the basis of their electronic structure. This classification is made by considering that there is some covalent character to the metal-halide bond, and that this covalent character can be described in terms of the occupancy of certain orbitals of the metal ion by electrons from the ligand. For coordination number six, one S , three P , and tq-o D orbitals are involved. If the orbitals are ( n - 1) 0 2 nS nP3, the complex is called an “inner orbital” complex; if the orbitals are nS nP3no2,the complex is called an “outer orbital” complex Reactions of inner orbital complexes, in general, are rapid if a t le& one of the three remaining (n - 1) D orbitals is not occupied by electrons from the metal. If these ( n - 1) D orbitals are occupied, the complex gives slow reactions in general. Several halo complexes of this last type, giving slox rates of hydrolysis and sloir- rates of substitution, are known. Probably the most familiar hydrolysis reaction falling into this category is the elon conversion of chromium(II1) from the green form in concentrated hydrochloric acid to the violet form in very dilute acid. Most complexes of the outer orbital type exhibit rapid reactions, but if the central atom of the complex is sufficiently electronegative, resulting in considerable covalent character to the metal-ligand bond, the substitution reactions become slon enough to be measurable. Most of the kinetic data t o be discussed are for an ion of this last type, SbCIe-. It is well to bear in mind that the mechanistic features displayed in the hydrolysis of SbCle- may not be characteristic of all halo complexes, and possibly may not even be characteristic of all outer orbital halo complexes. METHODS FOR MEASURMENT OF KIXETICS
A spectrophotometric method of following the kinetic has proved most successful in the case of S b C k .
Since all halo
complexes exhibit a “charge-.transfer“ absorption in the ultraviolet region, this method should be of general utility in studying the hydrolysis of these ions. Because these complexes are generally studied in the presence of the corresponding halogen acid the wave length region, which can be observed, is primarily determined by the absorption of the acid itself. For chloro complexes in hydrochloric acid the loxer wave length limit imposed by the solvent is about 210 mp, varying somewhat with the acid concentration and the instrument. I n general the molar absorbance index a t an absorption maximum is the order of 10,000, so that solutions as dilute as 10-411 in t,he complex can be used satisfactorily. The general characteristics of this absorption in the ultraviolet region can be seen from the several examples in the papers by Rabinowitch (10) and by Rogers and coworkers ( 2 , 6 ) . T h e spectral Characteristics of antimony(V) and its dependence on time are indicated in a previous article ( 7 ) . In actual rate studies the change in absorbance after dilution can be observed a t several selected wave lengths. The results of several experiments ( 9 ) of this type are discussed in the next section in connection n-it,h the The methods used to obtain the mechanism of the hydrol! desired rate constants from the experimental dat’a will noiv be indicated without going into detail. If the reaction goes to completion-Le., essentially all of the SbCl6- is converted t o one or another of the hydrolytic species-a plot of log ( D - D,) cs. time gives a straight line from whose slope the rate constant, for hydrolysis can be obtained. If the reaction goes only to an equilibrium [-where an appreciable fraction of antimony(T’) still remains as SbC16-1, the log ( D - D,) vs. time plot gives (I.h kf)>where k , is the pseudo-first-order rate constant for the formation of SbC1,- from the first hydrolytic species, Sh(OH)Cl,-. One can obt,ain kj, from the same experimental data by plotting log ( D - D p ) us. time, where D p is the absorbance the system iyould have if all the SbC16- had been converted to the first hydrolysis product. This plot is not a straight line, but from its initial slope one can obtain k h . I n such a reaction it is thus possible to obtain both kh and b,. Two other methods ( 1 1 ) have been used for studying this reaction, but they are not as convenient or satisfactory as the spectrophotometric method. The concentration of SbC16- during the hydrolJ-sis reaction can be determined by a polarographic method, because the SbCl6- ion is reversibly reduced, giving a well-defined diffusion current. The hydrolytic species do not interfere since they are reduced irreversibly at’ more negat,ive potentials. The second method utilizes the fact that SbC16can be extracted (R-ith an appropriate cation to form an ion pair) into organic solvents, whereas the hydrolytic forms are not extracted appreciably. The choice of cation and solvent is, of course, important; the protonated form of Rhodamine 13 as cation and benzene as solvent have been used for quantitative determination of SbCl6- ( 11 ) . 11-hether the last two methods could be used for other hydrokysis r,,avtions xi-ould require investigation of the properties of the halo ions and their hydrolysis products. These methods ivould not seem to have the general utility of t’he spectrophotometric method.
+
BIECHASISTIC FEATURES OF HYDROLYSIS OF SbCla-
The reaction is, as expected, first order in SbCls-. I n any given experiment t h e , concentration of all other reactants x a s far in excess of that of SbC1,-, hence each reaction appeared pseudofirst-order. The effect of other reactants \vas determined by observing t h e change in the pseudo-first-order rate const’ant, k,,, resulting from varying their concentrations from one experiment to another. Because the dependence on water remains undetermined, it is impossible t o say whether the water molecule enters the complex simultaneously with or subsequent t o the loss of the chloride ion, Or, in other words, it is impossible to say whether the reaction is of the SN-1 type or of the SS-2 type.
V O L U M E 2 7 , N O . 11, N O V E M B E R 1 9 5 5
1685
There is a stateIiient, in t,he earlier literature ( 1 4 ) ,based on the observation of formation of precipitate, that the rate of hydrolysis of SbC16- is decreased by increasing acidity. From equilibrium data (‘7) it is apparent t h a t formation of a precipitate is not a valid measurement of the hydrolysis of SbCI6-; in 6Jf hydrochloric acid SbC16- is about 99% hydrolyzed, but there is no formation of precipitate. And even a t acidit,ies \?-here all of the antimony(V) precipitates the rate of hydrolysis of SbC16- may not be the rate-determining step in the formation of the precipitate. Cheek (1) has shown that increasing acidity actually increases the rate of hydrolysis. His experiments TTere performed in hydrochloric acid of varying concentrations, so that [H+], [Cl-1, and ionic strength n-ere all varying. To observe the exact nature of this hydrogen ion dependence, experiments have n o r been performed (9) in solutions of constant [Cl-1, but varj-ing [ H +], by use of h~-drochloric acid-lithium chloride mixtures. At a total chloride concent’ration of 611 the rate constant ohserved is given by the expression k h = (3.9 0.8 [€I+]) X 10-3 niin.-l. At a total chloride concentration of 931 the expression min.-’ was found to hold. ki, = ( 5 . 3 1.6 [ H + ] )X
+
+-
3.6
I
I
I
I
I
1
I
I
the stronger the acid is the less likely it is to exist. Experiments n-ere tried n i t h tin(IT-), iron(III), and aluminum(III), but in none of these was there any effect on the hydrolysis that could be attributed to the corresponding chlorides. USE OF KINETIC DATA TO IDENTIFY SPECIES IN SOLUTION
Occasionally kinetic data give useful evidence about the ionic species in solution, although it rarely is conclusive unle ported by other evidence. Two examples of such a uti1 of kinetic data are given. The first example comes again from the antimony(1-) system, and concerns the identity of the first hydrolysis product. It has been stated that this product is Sb(OH)Clj-, without indicating any of the evidence. The kinetic evidence comes from measurements of the rates of hydrolysis and formation of SbCls- in solutions 9Jf in chloride ion, and of varying hydrogen ion concentration. In solutions of such high chloride concentration the hydrolysis, from the point of view of equilibrium, does not go beyond the first step. On t,he basis of the discussion of the types of hydrolytic species to he expected, it is reasonable t o assume that the first hydrolysis product is either Sb(0H)Clj- or Sb(H20)Clj. If the product is Sb(OH)Clj- then the rate of formation of SbC1,- is equal to k,[Sb(OH)C16-], and the rate of its hydrolysis is equal to kh[SbCI,-]. At equilibrium the tT7-o rates are equal, and [SbC16-]I[Sb(OH)Clj-] = kj/lih. Also a t equilibrium
k‘=
aE?OaSbC16-
.
~ R I ~ ( o H ) c I ~ -Uti-. ~H+
Although it is not completely cer-
tain that the use of lithium chloride-hydrochloric acid mixtures of constant chloride concentration guarantee a constancy of the activity coefficients, it is certainly the best t h a t can be done experimentallj- in this regard. Assuming that the coefficients are reasonabl!- constant
A similar analysis for the assumption that the first 111-drolysis product is Sb(H2O)Clj gives 1
I 2
I 3
I 4
I
I
I
I
5
6
7
8
9
[SbCls-] - k,l [Sb(H2O)CLI kh’ ~~
[H+l Figure 1.
=
kfflkh as a function of [Hi] for solutions 9 M in chloride ion
-,
and K’
=
[SI-,Cle-] [Sb(H20)Cbl
Figure 1 s h o w tbe experimentally observed dependence of on [H+]. The data can be seen to mtiafj- xell the relationship k,/kh = K’[H+]. The kinetic evidence then supports the position that Sb(0H)Clj- is the first hydrolytic species. -411 the equilibrium evidence ( 7 ) also supports this position. The second example of the use of kinetic evidence for identification purposes arises in regard to the chemistry of the synthetic element astatine ( 8 ) . The half life of this element is such that the chemistry of the element can only be studied on the tracer scale, and hence conclusive evidence in regard to oxidation state and ionic form is difficult to obtain. If a solution of astatine in concentrat,ed hydrochloric acid is treated Kith chlorine, the astatine is converted to a form that can be extracted into ethers. This behavior suggests that the astatine is present as either AtC16- or AtCI,-. (The ether extractability of both of these li-ould be expected because of their similarity to SbC16- and and IC14-.) If the ion is AtCls- it should display a sloTver rate of hj-droll-sis than SbCl6-, because of t,he greater electronegativity of astatine(V) compared t o antimony(V), and the rate could be measured by extraction experiments. ;ictually the extraction esperiments indicate a rapid reaction. On this basis it is unlikely that the astatine exists as AtCle-, but rather it appears t o be k,/lih
There are then two mechanisms operative under t,hesc conditions. The first is independent of hydrogen ion, and is the reaction of SbCl6- by either an SS-1 or SS-2 mechanism. T h e role of the hydrogen ion in the second, and hydrogen ion dependent, mechanism is undoubt,edly attack on the chloride. The data thus point to the existence of a species HSbCl6, but whether this is a true chloro-acid of reasonable lifetime, or is better described as merely a transition state, cannot be decided. It was also observed t h a t the presence of antimony(II1) accelerated the hydrolysis of SbCls-. T h e most, reasonable explanation of this effect, and one that is consistent n-ith the fact that the acceleration is greater a t lovier acidities, is that antimony trichloride is the effective form of antimony(II1). Its action, like t,hat of hydrogen ion, m-ould be in abstracting a chloride ion from t’he complex. From the magnitude of the effects observed antimony trichloride is several orders of magnitude more effective than hydrogen ion in this action. These results suggest that any Lewis arid should be effective in t,he same way. The difficulty n-ith aqueous sJ-stems is, of course, the fact that few Lewis acids exist as such in water, and
aH20aJbClsaSh(HzO)CIs a C l
1686
ANALYTICAL CHEMISTRY
AtCla-. In this regard the behavior of astatine is identical t o that of iodine, which forms IC11-under the same conditions. GENERALIZATIONS AND CORRELATIONS
For complexes of the outer orbital type, and which differ only in the central atom, slower rates correspond to increased covalent character in the metal-halide bond. There is ample evidence of this from substitution reactions in general ( I S ) , and this generalization has been used in the argument concerning the form of astatine in the last section. If Complexes of a given metal differing only in the halide ligand are considered, one might be inclined from the statement of the previous paragraph to feel that chloro complexes hydrolyze more rapidly than bromo complexes, because the latter are more covalent. It is more likely, however, that the rates of halide replacement in these inorganic systems parallel those in organic halides-that is, iodo complexes hydrolyze more rapidly than bromo complexes, which in turn hydrolyze more rapidly than chloro complexes, etc. I n these comparisons extent of covalent character is probably not the determining factor; polarizability and ion size are probably more important. It had been reported ( 7 ) that the hydrolysis of Sb(0H)Cls- is more rapid than that of SbCle-. Unfortunately no comparative rate data are available for Sb(OH)Cla- and Sb(OH)pCla-. If the trend continues in the direction indicated it would suggest that the remaining metal-halide bonds become more ionic as successive halide ions are removed, leading to the faster rates of hydrolysis. illthough the metal-oxygen bond is different from the metalhalide bond, similarities may be expected between the reactions of inorganic oxy anions and the halo anions. The reaction in the oxygen system corresponding to the hydrolysis of the halide
8th Annoal Summer Symposium-Role
system is the oxygen exchange between water and the oxy anion. iicceleration of the oxygen exchange by hydrogen ion has been observed for many oxy anions (3, 6 ) , and mechanistically the effect of hydrogen ion can be visualized to be the same in the two cases. There are several examples of halide substitution being accelerated by Lewis acids, similar to the influence of antimony trichloride on the hydrolysis of SbCls-. I n the organic field Lewis acids, most commonly .4g+ and mercuric chloride, are often used to accelerate the hydrolysis of alkyl halides ( 4 ) . An example in inorganic systems is the catalytic effect of platinum(II1) on the radioactive exchange between C1- and ’PtC&-(12). There ii; still much to be learned about these reactions. LITERATURE CITED
Cheek, C. H., Ph.D. thesis, Washington University, St. Louis, N o . , January 1953. DeSesa, hI. A . , and Rogers, L. B., Anal. Chim. Acta, 6 , 534 (1962).
Hall, E.F., and Alexander, 0. R., J . A m . Chem. Soc., 62, 3455 (1940).
Hughes, E. D., Quart. f2ei.s. (Lor(dorij, V, 245 (1951). JIerritt, C., Hershenson, H. II.,and Rogers, L. B., ANAL. CHEN.,25, 572 (1953). llills, G . A , , J . Am. Chem. Soc.. 62, 2833 11940). Seuniann, €1. AI., Ibid., 76, 2611 (1954). Keuniann, H. h l . , unpuhlirhed data. Keurnann. H. II., and Haniette, K. IT., unpublished data. Habinowitch, E., Recs. M o d . Ph//s., 14, 112 (1942). Ramette, R.IT., Ph.D. theais, University of Minnesota, Minneapolis, Rlinn., June 1954. Rich, R. L., and Taube. H., J . -Am. Chern. Soc., 76, 2608 (1954). Taube, Henry, Chem. Revs., 5 0 , 09 (1952). Weinland. R . F.,and Srhmid, H . , Z. anorg. Chem.,44, 37 (1903). RECEIVED for rerieiv J u l y 2 2 , 1955. .%ccei,ted September 6, 1955.
of Reaotion Rates
Competing Rates of Oxime Formation Determination of Aromatic Aldehydes in Presence of Aromatic Ketones LEWIS FOWLER Monsanto Chemical Co., St. Louis, M o .
In the analysis of mixtures of organic compounds containing a common functiona 1 groupTmethods based on differences in rates of reaction with a common reagent are particularly useful. The example of the rates of oxime formation with mixed aromatic aldehydes and ketones, specifically vanillin and acetovanillone, illustrates the general technique. Calibration curves relate the concentration at a chosen time with the original concentration. The calculation of the extent of reaction with time is illustrated. The selection of the optimum conditions of reaction is desciibed.
THE
4 usual methods of analysis applied to mixtures of two constituents which differ only slightly in chemical properties often lead to unsatisfactory results. This situation occurs frequently in organic syntheses where two homologous products having the same functional group are obtained. It is not always practicable t o utilize a physical method for quantitative analysis of such mixtures. A technique which can sometimes solve these problems depends on the fact that minor differences in substituents or structure frequently cause pronounced changes in the relative rates of
reaction with a given reagent Kolthoff and Lee (6-7) cite examples of the use of this technique for the analysis of mixtures of esters, aldehydes, ethylenic compounds, etc. Hass and Weber ( 4 ) analyzed mixtures of primari isoamyl chlorides by the rate of reaction v i t h potassium iodide. A method based on the application of rates of oxime formation enables vanillin to be determined simply in the presence of acetovanillone (4’-hydroxy-3’-methoxyacetophenone)with excellent precision and accuracy ( 2 ) . p-Hydroxybenzaldehyde and syringaldehyde behave quantitatively like vanillin, and acetobehaves syringone (4’-hydroxy-3’,5’-dimethoxyacetophenone) like acetovanillone. Trials of the method (without attempting t o determine optimum conditions) Tvere made on several aromatic aldehyde-ketone pairs with favorable results (Table I). The principles underlying this example illustrate the general criteria for using competitive reaction rates as an analytical tool. The reaction of carbonyl compounds with hydroxylamine to form oxime is second order (3). There is general acid catalysis, but the use of half-neutralized hydroxylamine hydrochloride in excess buffers the system sufficiently to make this effect negligible (IO). The relative rates for a number of aromatic aldehydes and ketones have been compared by Vavon and Montheard
