Titration Characteristics of Organic Bases in Nitromethane CARL A. STREULI Central Research Division, American Cyanamid Co., Stamford, Conn. k Nitromethane is an excellent solvent for the titration of weak organic bases. Amines, amides, and ureas, however, show diverse titration characteristics in this soivent. This study was undertaken to correlate structure with titration behavior. Amines show normal titration behavior in this solvent but monosubstituted and unsubstituted amides and ureas show extremely steep titration curves. The behavior is probably due to hydrogen bonding between solute molecules. In nitromethane amides, ureas, heterocyclics, and hydroxyamines show relatively greater basicity than they do in water. Equations are presented relating basicity of the various classes of compounds in nitromethane and in water.
N
as well as a c e tonitrile, is an eminently practical medium for the titration of weak organic bases. A very weak acid, nitromethane has an extensive potential range. neither levels nor reacts with most solutes, and has a dielectric constant of nearly 40, which allows potentiometric measurements to be made easily. The insolubility of most salts in nitromethane and the effect of \vater on the potential range are its principal disadvantages. Traces of water in the solvent make impractical the quantitative determination of bases ~ i t hpK, values greater than 14. Fritz and Fulda (1) ha;re shown the uses of nitromethane as a differential medium, but the solvent mixture they used contained 20% of acetic anhydride. Hall (3) investigated the behavior of amines in several nonprotolytic organic solvents and concluded that to a first approximation pK.(H20) values are a good index of the base strength of these compounds in other solvents. His principle investigations were carried out in acetonitrile and ethyl acetate, but he also studied behavior in ethylene dichloride, nitrobenzene, and nitromethm e . Other studies have been made for dcetic acid (/t) and acetic anhydride (9). %cause of the utility of nitromethane :rid the limited data for this solvent, the work described was initiated. The compounds selected for use in this work were those which offered a variety of ITROMETHANE,
1652
c
ANALYTICAL CHEMISTRY
molecular organic structure and pK. range consistent with basic properties. Primary, secondary, and tertiary amines, heterocyclics, ureas, amides, guanidine derivatives, and imidates have been included. EXPERIMENTAL
The nitromethane was obtained from the Fisher Scientific Co. It was reagent grade material and was used without further purification. Standard 0.05N perchloric acid solutions were prepared by diluting 4.2 ml. of 72% acid to 1 liter with nitromethane. Solutions were stored in closed, brown bottles, and were standardized against potassium hydrogen phthalate dissolved in acetic acid ( 7 ) a t weekly intervals. The standard solution appears to be stable for a t least a month. Most of the materials titrated were Eastman Kodak JWite Label grade. The cyanoethy! compounds and imidates were research samples prepared and purified in these laboratories. Apparatus. All titrations were perReagents.
formed using a Precision-Dow Recordomatic titrator. Glass and aqueous sleeve calomel electrodes were employed to measure potentials during titrations. Solutions were agitated with a magnetic stirrer. Procedure. A millimole of material was dissolved in nitromethane and diluted t o 100.0 ml. in a volumetric flask. An aliquot of 25.0 ml. of this solution was then diluted t o 100 ml. with nitromethane and titrated. Three aliquots were run on each compound and blanks were determined on each batch of nitromethane. RESULTS AND DISCUSSION
All titration curves were corrected for blanks and plotted a8 potentini against per cent neutralization. Several examples of the plots are shown in Figure 1. The titration curves illustrated include examples of amines, ureas, and amides. No leveling of base strength was observed even for amines as strong as piperidine. Because nitromethane does
Figure 1. Titration of nitrogen bases in nitromethane a. b.
Diphenylamine N-Methylacetanilide E. Urea d. N-Ethyl-N-methylaniline e. Pyridine f. N,N-Diethylaniline g. N,N'-Diphenylguanidine
d
-PE
E
% Nsutralfzrd
600
700
i
-I
600
500
400 I
,i i
300
/'
-2001 -2
'
I
0
'
2
'
'
4
'
'
6
'
' ' '
8
IO
'
100
;
I
I
40
a0
PKa (H;P)
Figure 3. u.
b.
.-t
I
I
I
m
00
06 NeutraliPcd
Figure 2. Basic strength of nitrogen compounds in nitromethane and in water
have a labile hydrogen, leveling probably does occur, b u t only in a more basic region. All the amines, heterocyclics, and guanidine derivatives titrated gave curves qualitatively similar t o those illustrated for these types of ..ompounds. Between 20 and 80% iieutraliaation, the titration curve for ziiose bases with pK. (HpO) values less than 8 are essentially linear with a slope of 1.1 f 0.1 mv. per per cent :ieutralization. Tt can also be seen ehat pyridine, which in water is a weaker base than either N,N-diethylmiiine or N-eth>~l-S-meth!;laniline, s h o w a base strength greater than the iatter compound when dissoived in i;itromethane. Urea is also consider2%:; stronger than diphenylamine in this soivent, which is contrary to literature pi(, data (4, 7). The shape or &he titration curve for wea is interesting, being much steeper *.. .&an Che curve for smines. The slope between 20 and SOYG ~ieutralization is rl u , mv. ~ per per wiltj neutralization. This type of curve is characteristic of si: the ureas as ivd; as mono and mubstituted amides. Disubstituted .inides show titration i'urves similar to :hose for amices. This type of Ge,iavior was : l a noted in either acetic teid (4) or acetic anhydridc (9),but is characteristic for phenols in both acetone ( 2 ) and pyridine (IO). It is probably due to intermoiecular bpdrogen bonding between amide or urea
I
50
c.
Titration of nitrogen bases in nitromethane
Acrylamide 2-Pyrroiidinone Trir-2-cyanoerhylamine
molecules rather than hydrogen bonding between solute and soivent. Disubstituted amides with no svailable hydrogen atoms behave like the amines as exemplified by N-methyj acetanilide (Figure 1). Results for ail the compcunds titrated with the exception of some ~ e r y weak amides were quantitative. i:i the latter cases the potential change is so slow that it is difficult t o lorate an end point. Figure 2 illustrates the relation between pK. (HzO) and A " P (CRs KO2) for the compounds titrated. Pertinent data are given in Table I. h"P values were calculated bj. algebraically subtracting the nalf neutralization potential ( H N P ) for a N,N'-diphenylguanidine sample run on the same day from the H N P 7.alue of the compound being tested. This is necessary because 3 f d:iy-to-day shifts in liquid junction potentials of the electrodes. a " P vaiues are reproducible within 5 to 6 rniilivoits, while the EiNP value of a compound using this electrode system may shift i00 mv. over several days' time. N,N'diphenylguanidine was chosen as a reference standard because of its high base strength, purity, and availability. :'sing this compound as a reference, ail titration curves and G I N P values
o. Ethyl p-nitrobenzimidate a, Ethyl benzimidote
stand in c6rrect relative relation to one mother. Two iinear relations between pK, (H20) values and m N P (CH3NO2) wlues are apparent in Figure 2. 3 e main sequence is determined by the amines, the shorter sequence by amides and iweas. Heterocyclic nitrogen compounds and hydroqlated amines do not fit either sequence. Essentiaiiy h i s jl!ustrates that compounds ron.\, f l 3 /' tiiining the grouping ,S--N, as nclI 2s heterocyciics and hydroxy smines are, ;elati?:e Lo amines, stronger bases In ;iit;ornei'nane :han they are in water. e . i keiiavior was not notcu e? .iliiLureas, in either acetic :.'r scesic anhydride (9) b u t .:aids ;rue for such solvents 3s arzto;iitriie or acetone. it is again most >rc.hhiy ;elated to hydrogen LondiKg bc:wcen sclute and solvent moiecules i i l ::le various solvents. TLe straigrlt m e drawn for the amine sequence 1s ihe :easr, ncuares solution of ;he data for nil t h e aniines exc;usi.vre of the heterocyciic anti :.iy(:roxvamines. T!:e equation ieiating ::i< i&O) ,, and S i Y P (CH,NO,) values ror the emine 3enuence is ?KO (&O) = io.12 - 0.0129 x AHh'P (CHaYOz) (1) VOL. 31, NO. TO, OCTOBER 1959
0
*553
Table I .
Basicity Data for Amines, Amides, and Ureas in Nitromethane and Water
Compound
AHNP
.V,S’-Diphenylguanidin- the broken line in Figure 2 . Calculated pli, (H20)valuzs using a ” P (CH3S02jdata and the appropriate equation are listed in the third column of Table I ; deviations between literature and caiculated values are listed in column 4. The standard deviation for all the data is 0.27 pfi, unit exclusive of the heterocyclics and hydroxy amines. Both equations fail badly for these conipounds which show
1654
ANALYTICAL CHEMISTRY
I ? 09 10 64
10 12 10 72
10 67
11 01 10 60
9 02
9 12
86 6 52 5 99
7 49
I
5 5 4 4
21
14
52
11
6 43 6 20 5 a0 5 14 1 81 3 80
0 12 -0 3: 0 3; -0 07 (’ 10 -0 3‘: -0 09 0 21 0 00 0 00 0 ‘I-0 31
CONCLUSIONS
602
Bis-2-hydroxyethylaniine Tris-2-hydroxyethylaminc
Compound 2-Pyrrolidinone .V-Methyl-2-pyrrolidinone .icrylamide Xitrilotrispropionitrile Triphenylphosphine Ethylbenzimidate Ethylacetimidate E thylcarbethoxyacetimidate Trichloroacetimidatc, p-Xitrobenzimidate Ethyl a-chloropropionimidate
PK, (Lit.) 10 00
5.06
i 6 191
5.30 8 87
(10 07’1
7.82
0.90
0.50 0:61 -0.60 -0.50
-0.48 -0.30
-1.70
i6 2c)
19 49)
0 79 0.61 0 4s -0.17 -0 34 -0 70 -0 7 0 - 1 56
i 13 0 9F: l.2G 1 6;. -0 1 1 0 li -0 1 : ; -0 I S -0 le, c) 22 c1 46
-0
11
It is possible to resolvr, mixtures of amines in the mais srqucnce which differ in I K S F vslues by approximately 180 m x . The theoretical limit of resolution should be about 100 mv. Honewr. this has no: been achieved with any of the compounds tested.
is, (6; (11) ( .!1>’,
(4’1 ( +;, (6)
(4) (4; (4)
(4;
(7:
strengths interinediate between the two main classes of compounds. The two lines are essentially parallel with slopes of 78 mv. per pK. unit. Hall ( 4 ) assumed a slope of 59 mv. per pK, unit for this relation in acetic acid. In acetic anhydride this value is 51 mv. per pK. unit (9). Nitromethane should, on this basis, have greater resolving power than the anhydride. A number of other molecular species with basic characteristics b u t unknown or poorly defined pK, values were also titrated. Typical titration curves are given in Figure 3 and AHNP (CH,NO,) values in Table 11. pK, values have been calculated from these tiatn using either Equation 1 or 2 depending on the molecular structure of the solute. I n the cases of the imidates and thc phosphine, Equation 1 \\as used since the titration behavior of these compounds was similar t o that of amines. Assuming that this equation applies in these cases, the 90% confidence limits for these. p& values are 1 0 . 4 pK, un;t. Kitromethane makes a convenient titration medium for imidates which are hydrolyzed in water to esters Very weak bases such as dimethylsulfoxide. tetrahydrofuran, and acetonitrile show no titratable basicity in this solvent. This is most probably due t o interference by water in the titrant.
The foregoing data, with the studies made in acetic acid ,.$I and acetic anhydride (9),illustrate that while the relative basicity n-ithiu class of compounds is not grenti?. affected by changes in titratinn medium, basicity relations between various classes such a i amides and amines are affected to a considerable deprre. This change appears to be related t c the hydrogen bonding chsracteristics o! the solvent. Sr. unhnown pIL, VZIUI= may be caiculatttl from nitro:nethane data if the class of the curnplsund is consiaered. pK, values so obtained are probabl! reliable within U.5 unit. Resolution of mixtures 3f compounds can be predicted on the basis of pK, data or nitromethane data if the shape of the titration curve and structure of the solute are considfvd. Mixtures of amines are more readily resolved than mixture? of amides or ureas because of the shape of the titration curvc. ACKNOWLEDGMENT
The iniidates used in this w r k $yere synthesized by F. C. Schaefer and Grace Peters of the American Cyanamid Co. The author thanks Seymour Sandler, who performed a great many of the titrations. LITERATURE CITED
Fritz. J. S.. Fulda. 31. 0.. ANAL. C H E M ; i i , 183f (1953).’ ( 2 ) Fritz, J. S.,\-amamur:i, J. J., Zbid., 29,1079 (1957). 13) Hall. H. K.. Jr.. J . Phus. Chem. 60. f 1) . - I
- - I
~
~
(4) &ll, ?;. F., J. Am. Chem. SOC.52, 5115(1930). (5) Iiolthoff. 1. M.,Furman, N. H., “Potentiometric Titrations.” p. 329, Riley, Ten- York, 1926. ( 6 ) Lange, K. A., “HandSook of Chemistry,” 4th ed., pp. 1 2 2 G i , Handbook Publishers. Sanduskv. Ohio. -,-1941. - ~ ( 7 ) LeMariej H., Lu&s, H. J., J . Am. Chem. SOC.73, 519s (195:). (8) Seaman, IT., Wen, E.. ANAL.Cmnr. 23, 592 (lj51). (9) Streuli. -. A,, Ibid.. 30,997 (1958). (10) Streuli, C. A, Kliror,, IC. R., Zbid., 30,1978 (1958). (11) Streuli, C. rl., Sandier, S., unpublished data RECEIVEDfor review hIxch 12, 1959. Accepted June 24, 1959.
