ALDENJ. DEYRUP
GO
titrated with 0.01 m ceric sulfate. T h e color change is from a faintly reddish-yellow to a clear yellow. I n the presence of large amounts of iron, as in steel analysis, the addition of phosphoric acid t o discharge the iron color becomes necessary. The indicator blank amounts to about 0.15cc. of 0.0t m ceric sulfate in a volume of 200 cc. of vz sulfuric acid. Table 111 contains the results of analyses which show the precision attainable on mixtures comparable to those obtained by solution of commercial steels. U'e have further analyzcd the Bureau of Standards chrome-vanadium steel KO. X c. -4 2-g. sample was dissolved in 20 cc. of 5 m sulfuric: acid, and after addition of 5 cc. of 8 m nitric acid, the sol.ltion w w evaporated to fumes. After dilution and filtration the insoluhle residue was ignited, treated with hydrof?uoric and sulfuric acids, and fused with sodium carbonxtc. The dissolved melt was added to the main solution, and vanadium was then determined by the method given for the determination of small amounts. The re,;ults are V (fourid),
70 0.2373
Dctn. 2
Av. dev. pts. per 1000 1. (certificate), yG 2.5 0 , 2 3 5 (range : 0.225-0.246)
[ C O N TKIDOTION P R O M THE
Vol. 5G
Two determinations in which the recovery of vanadium from the silica residue was omitted led t o values of 0.2349 and 0.2343, so t h a t the amount of vanadium thus carried down is easily detected by this method. It is worthy of note t h a t it is now possible to determine the vanadium content of a steel which carries only 0.2% of t h a t element with a precision of about 3 parts per thousand.
Summary Oxidation potentials relative to cach other of the ferric-ferrous system, of the phenanthrolineferrous indicator, and two stages of the vanadate reduction have been measured in 1, 3 and 5 m sulfuric acid. From these the necessary conditions for the separate oxidimetric titration of iron and of vanadium have been predicted. A rapid and precise method for the determination of vanadium in the presence of iron, chromium and molybdenum has been developed and tested. RECEIVEDSEPTEMBER 5, 1933
NEW YORK,N . Y .
LABORATORY O F PHYSICAL CHEMISTRY O F
THE
UNIVERSITY
OF \ v l S C O N S I N ]
Determination of Acidity in Ethyl Alcohol by Velocity of Acetal Formation BY ALDENJ. D E Y R U P ~ Thc: development of the theories of salt effects and general acid catalysis, due mainly to Bronsted, has placed reaction velocity measurement among the accurate methods of determination of acidity in aqueous solution.2s3 A satisfactory application of the method to non-aqueous solutions depends on the selection of a reaction sufficiently sensitive to hydrogen ion to permit of measurements a t low ionic strength. Adkins and Rroderick" measured the rate of such a reaction, the formation of diethyl acetal from acetaldehyde in alcoholic solution with hydrogen chloride as catalyst. The decrease in velocity constants as the reaction proceeded was attributed to t h e effect of water produced in the reacti0n.j This reaction has been investigated as a method of measurement of acidity in absolute ethyl alcoholic solutions. It was found convenient to follow the course of the reaction dilatometrically, the high sensitivity of this method making possible the use of dilute solutions of substrate. (1) ('2) (3) (4) (5)
S a t i o n a l Research Fellow in Chemistry. Hrtiiisted a n d Grove, THISJOURNAL, 51, 1394 (1!)30) Kilpatrick and Chase, i b i d . , 63, 173%(1931). Adkini aticl Bro(lerick, ibid.,SO, 17X 11928) 1:e I.ceuw d? p h y s i k . C h e i ~ .77, , 2 x 4 (1911).
Materials and Procedure Ethyl Alcohol.-Commercial 95% ethanol was dehydrated by refluxing over several portions of calcium oxide and distilling in an apparatus designed t o prevent access of moisture from the air. It was kept in a flask protected with calcium chloride tubes, and its water content was estimated a t intervals from density determinations a t 25'. The lowest density observed was 0.78505 g./ml., which is in good agreement with the density 0.78506 of Osborne, McKelvy and Bearcc.6 The density was found t o vary linearly with the water content from 0 t o 1%. The water content was therefore estimated froni the density figures for 99 and 100% alcohol in the "International Critical Tables." It was not allowed to cxceed 0.020/,. No removal of aldehyde was necessary since the acetal formation is first order, and aldehyde increases the density of alcohol as does water. Acetaldehyde.-Commercial acetaldehyde was distilled through a column of calcium chloride, the fraction boiling above 22" being rejected. HCl, HBr, HI, HC104.-C. P. concentrated acids were used. Since the concentrations of these acids as catalysts did not exceed 0.001 molar, the amount of water introduced with the acid is negligible. Picric acid, 2,4-dinitrophenol and oxalic acid were recrystallized from water, the last being dehydrated a t (6) Oshorne, LIcKelvg and 327 (1913).
Bearce, Bur. . S / a m i a r h B u l l . ,
'
Jan., 10N
DETERMINATION OF ACIDITYIN ETHYL ALCOHOL
100 '. Trichloroacetic and monochloroacetic acids were distilled in vucuo. Aniline, o-toluidine, dimethylaniline, and pyridine were purified by distillation. Pyridine was dried over potassiurn hydroxide before distillation, and separate samples wcrc titrated i n a sinall volume with standard hydrochloric acid in the presence of thymol blue and of brom phenol blue ai indicators. The mean result was 99,9yo pyridinc. Procedure.--For fast reactions the dilatometer comprised a niising bulb coiiiiected by a well ground stopcock t o a reaction bulb of about 40 cc. capacity. -4graduated capillary of about 0.4 sy. inm. cross section was sealed t o the reaction bulh. T h e scale and capillary were calibrated by the mercury thread method. T h e leakage through the stopcock, which was lubricated with a small amount of vaselinc, amounti?d t o approximately 0.2 mm. on the scale per hour. The same stopcock with a mercury seal was found t o leak more rapidly, presumably because the mercury dc'es not immediately force the solution (out of t h e interstices of the stopcock. For this reason a mercury seal was not used for the major portion of the work. T h e inhibition hy mercury of the hydrogen chloridecatalyzed reaction noted by Adkins and Broderick4 was not observed, however, in a series of experiments in which mercury in thc dilatometer bulb was used t o seal the stopcock. For slow reactions a dilatometer was designed without stopcocks t o eliminate leakage. To a 30-cc. reaction bulb were sealed two vertical tubes, one of which was a graduated capillary of 0.5 sq. mm. cross section; the other was a 2-nim. tube constricted t o about 0.5 sq. mm. cross section and bent downward just above the constriction t o ininiinize drainage. 'The bulb was filled through the constricted tube, anci readings were taken o n t h e graduated capillary after ad.justing the level in the filling tube t o a mark 011 the consxicted portion The reaction was started by mixing a sinall defiiiite voluiiic of a t' lygsolution of acetaldehyde in absolute alcohol with the solution containing the catalyst (immediately after preparation of the latter) and transferring t o the dilatometer by air pressure. Both solutions had previously been tarought t o temperature in a thermostat regulating to *0.~)01",and within five minutes after mixing the first rc:iding was takcii. The exact interval was not determined, since the reactions are of the first order. Tinic nicasurciiie its wcre niade with a Calibrated stop watch. It was not ftea:,ihle t o add pure aldehydc directly because of the large heat of solution in alcohol. Since the water produced inhibits the reaction the concentration of aldehyde was kept below 0.02 molar. T h a t this coiicentratiori affects neither the solvent properties nor the acidity was proved by the fact t h a t variation of thc initial concentration from 0.005to 0.1 did not affect the velocity coiista~it. T h c conclusion of de Leeuw5 t h a t t h e density of alcohol-acetaldehyde solutions increases with time was riot corrolm-atcd. When acetaldehyde and ethyl alcohol wcrc inixcd a t a temperature sufficiently below t h a t of the thermostat t o balance the heat of mixing, 110 volumc changc. wai observed from five minutes to three hour.; after mixing.
!j 1
The first order reaction velocity constants were obtained graphically, and were reproducible to +3% on separate runs. In Table I is part of a typical experiment with hydrochloric acid as catalyst (20 to 50 readings were taken usually). 1 represents the difference between the observed dilatometer reading, hit and that calculated from the graphically determined constant k and the initial and final readings, 11, and hi, by the equation kt
11
- h, h,
= 111 --
ht
-
1111 determinations were made a t 23.00", T.4BLE 1 DATAFOR TYPICAL REACTION HCl U.000217~2k(graph.) 0.0318 mill.1.
min.
0 1 2 4
8 12 16 20 32 40 50 (200)
h , (obs.), cm.
7.25 (h,) 7.18 7.09 ti. 93 ii. 6 4 6.40 ti. 18 3.99
hl (calcd.), cm.
... . 7.17 7.08 0.93 6.66 6.41 6.19
5.57 3.37 5.17
G.OO 5.56 5.34 5.14
4.GO(h/)
. ...
1, cm.
....
-1r.01 - ,01 I!(!
02 .01 11 1 01 - .ill -
. [I3
- ,(I3 .
, .
Results Strong Acids. Catalytic Constant of Hydrogen Ion.-The catalytic constants k C (where k is velocity constant in min.-l and C is concentration of acid in moles per liter) of HC1, HBr, H I and HClOA were determined in solutions from 0.00003 to 0.0005 molar. The mean values were. HC1, 143; HBr, 112; HI, 113, EIClO,, 141. They were independent of concentration and unaffected by addition of 0.001 molar salt with a common ion (LiC1, XaBr, KI, SaC10,i. These facts confirm the conclusion of NurraxRust and Hartley' that these acids are completely ionized in dilute alcoholic solution The contrary conclusion of Schreiner' with respect to HC1 and HBr depends on catalvtic esterification in relatively concentrated solutions of electrolytes. The average xalue 143 ma\therefore be taken as the catalytic constant, kH O H ) , of solvated hydrogen ion in absolute alcohol a t 2-5O for low concentrations of electrolyte ( i ) \lurra)-Iiust ani1 IHartlev I'roi K u i > o Al26, 14 lll?rO h 5 c h r e i n c r 7 p h i i k Ciieiii 111 41'l 1'124
ALDENJ. DEYRUP
62
The temperature coefficient and energy of activation for the reaction catalyzed by solvated hydrogen ions were obtained from the mean catalytic constant, 494, for HCI a t 3.5'. E,,,.
=
22,600 cal.
The reaction shows a linear primary salt effect (Fig. I ) , the same for the different 1:1 electrolytes 300
1
io01 (1
I
I 0.02
Fig. 1.-Primary
I
I
I
I
I
I
VOl. 56
stant. It may be concluded therefore that the reaction is not catalyzed by undissociated nitric and picric acids, and that it is not catalyzed by acids weaker than H I(C2HbOH) in alcoholic solution. The PIC (negative logarithm of the dissociation constant IC,) of nitric acid is plotted in Fig. 2 against the square root of the ionic strength. The slope, 5.9, of the experimental curve up to db = 0.06 is approximately that calculated from the Debye-HLickel limiting law, 5.6. The variation of PK with dii for picric acid determined by Larsson'O (conductivity) and by Gross and Goldstern" (spectrophotometric) are shown in the same figure for comparison. The slopes are, respectively, 7.6 and 7.7.
1
0.04 0.06 0.08 0.10 Ionic strength, p. salt effect. 0, NaC104; 8 ,LiCI; 0 , K I ; 9 , NaI.
LiCl, KI, XaI, NaC104 up to ionic strength 0.1. The accuracy of the data from which this conclusion is drawn is probably not better than :k 10% a t the higher electrolyte concentrations since it depends on measurements in unbuffered solutions. The salts used, however, were carefully recrystallized. It is to be noted that this salt effect is considerably greater than known primary salt effects in aqueous solution.9 It ma!. be expressed by the equation p =
~ H + ( c ~ H . o H= )
143
+ 1500 p
For a measure of acidity in solutions containing more than a few thousandths molar electrolyte it is necessary to correct the catalytic constant by the above equation. This has been done where necessary in the work here reported. Weak Acids. Change of Electrolyte Concentration.-The dissociation constants of some weak acids were determined by measurement of catalytic activity of their alcoholic solutions alone or with their salts. No measurable change in the velocity constant was obtained for nitric and picric acid buffers when the buffer concentration was changed from 0.0005 to 0.005 molar, keeping the buffer ratio and ionic strength con(9) Br'insted and Grove (Ref. 2) and Kilpatrick a n d Chase ( R e f . 3) found a n unusually large primary salt effect on t h e reverse r c x t i o n :.kceta1 hydrolysis) in aqueous solution.
4.0 0
0.02
0.04
0.06
0.08
J
0.10
,'t
Fig. 2.+K of weak acids as function of electrolyte concentration: Curve 1, nitric acid a t 25"; added electrolyte: 0 , none; C3, sodium nitrate; 0 , silver nitrate. Curve 2, theoretical slope calculated from Debye-Huckel limiting law. Curve 3, picric acid a t 20' (Gross and Goldstern). Curve 4, picric acid a t 25 (Larsson).
In Table 11, column 2 , are given the mean calculated dissociation constants of three carboxylic acids, two phenolic acids, and nitric acid. For intercomparison they have been extrapolated to zero ionic strength by the equation log K O = log KO + 5.9 d k As a check the dissociation constants of picric acid and 2,4-dinitrophenol were determined colorimetrically, using a Bausch and Lomb colorimeter, and are given in the table in paren(10) Larsson. Thesis, Kopenhagen, 1924. (11) Gross and Goldstern, Monalsh., 66, 316 (1930).
DETERMINATION OF ACIDITYIN ETHYLALCOHOL
Jan., 1034
theses. The discrepancy in the case of 2,4dinitrophenol is probably due t o the necessity of determinations in unbuffered solutions with this weak acid. For the same reason the dissociation constant of monochloroacetic acid is not as precise as those of the less weak acids. T A B L E 11 DISSOCIA~ICIN CONSTANTS O F
KO Acid
Nitric Picric Trichloroacetic Oxalic Monochloroacetic 2,4-Dinitrophenoli
E t . alc.
2.7 x 1.0 x (1.0 x 3.5 X 2.6 x 1.8 X 1.8X (4.2 X
io--* 10-4 10-4)
lo-' 10F lo-*)
WEAK
ph'u
ACIDS PK
E t . alc. W a t e r AgK
... . . .
3.57 4.0
0.8 3.2
5.46 6.58 7.74 7.74 (7.38)
.7 4 . 8 1.3 5.3 2.9 4 . 8 3.9 3.8 (3.5)
The dissociation constant found for nitric acid agrees in order of magnitude with the figure 0.9 X l o T 4 of Murray-Rust and Hartley' from conductivity measurements. The constant for picric acid is in good agreement with the constants of Larsson'O (0.8 X lo-*) and of Gross and Goldstern" (1.8 X lop4, 20'). In columns 3 and 4 are given the values of PI< in alcohol and water,I2 respectively. A PI- t h a n phenolphtha!ein as indicator J t i color change correspond, more c l n x l y with t h e 1'31 of stoichiometrically neutral ,?diiim sulfite solt,tii)ns, t h u s making unnecessary t h e use of cornpariaon solutions of i n d i m t o r .
Summary I . The rate of formation of diethyl acetal is suitable as a measure of acidity in absolute alcoholic solution at relatively high or low ionic strength. The reaction is not catalyzed by acids weaker than hydrogen ion, and a correction may be applied for the large primary salt effect. 2. Hydrochloric, hydrobromic, hydriodic and perchloric acids are completely ionized strong acids. Nitric and picric acids are weak acids of about the samc strength in this solvent. 3 . The effect of change of electrolyte concentration and solvent on the dissociation constants of a number of weak acids and bases has been determined. 1. The inhibition of the reaction by traces of water may be ascribed to its basic character. The base strength of water in alcohol has been estimated. 5. Catalysis of the reaction in the presence of calcium chloride, calcium nitrate and lithium chloride is due to traces of strong acids. AIADISOS, \YISCOSSIX
RECEIVED Sb PTEMBBR 7 , 1933
