in determining reactive alkyl halides have enabled alkoxy1 determinations on higher alkyl groups which are not amenable to the Zeisel method. For example, the benqloxy group of carbobenzyloxy compounds, benzyl ethers, and benzyl esters is quantitatively converted to the benzyl bromide which, after extraction with toluene, was determined by Procedure 11’. This continuing.
LITERATURE CITED
(1) Bender, M. L., J. Am. Chem. SOC. 75,5986 (1953). (2) Fritz, s. J.1 Lisicki, N. b1.1 A S A L . CHEM.23,589 (1951). (3) Katchalski, Ephraim, Advances i n Protein Chem. 6, 123 (1951). (4) Patchornik, Abraham, Ph.D. disserta-
tion, Hebrew University, Jerusalem, Israel, 1956. ( 5 ) Patchornik, Abraham, Berger, Arieh, Katchalski, Ephraim, J . Am. Chem. SOC.
79,5227 (1957). (6) Ibid.,. 79, 6416 (1967). (7) Riddick, J. A., ANAL.CHEX 30, 793 (1958). (8) Vogel, A. I., “Practical Organic Chemistry,” p. 175, Longmans, Green, London, 1948.
RECEIVED for review October 16, 1958. Accepted February 16, 1959. Work suported b research grant (PHS H-2279) From U. . National Institutes of Health, Public Health Service.
8
Conductometric Titration of Very Weak Acids FRANC0 GASLlNl and LUClO ZION NAHUM Research Division, Vita Mayer &
Co.,Milan, Italy
,The determination of weak and very weak acidic groups, particularly phenols, is carried out by dissolving the sample in an aqueous solution of a weak nitrogenous base (ammonia) present in excess and performing the titration with lithium hydroxide conductometrically. Mono- and polybasic acids (carboxylic and phenolic acids, enols, imides) have been titrated in this manner with accurate reproducible results. Intersection angles are obtained which are as satisfactory as those given by strong acids using the usual conductometric method. The method makes it possible to reveal and titrate very weak acidic groups which were not revealed by the usual conductometric method. A better differentiation i s obtained between different functions of polybasic acids.
T
conductometric titration of very weak acids with strong bases ( 1 ) is hindered by hydrolysis phenomena which, as the neutralization proceeds, cause a progressive nonlinear increase of the conductance of the solution, due to the increase of the hydroxyl ion concentration. Therefore, the weaker the acid to be titrated, the more inclined and shorter the straight line with which the titration curve begins. ilsa consequence, large angles are obtained between the neutralization line and the excess base line, which are a source of inaccuracy in the location of the end point; also the accurate extrapolation of the initial straight line becomes, as a result of its brevity, difficult or even impossible. hforeover, in the case of weak acids in very dilute solution it is often impossible to get quantitative results-e.g., nonquantitative ralues are obtained in the titration of phenols in water-alcohol solution with lithium hydroxide, by KoltHE
tion equilibrium of the acid, HA: HA B aABH+ If the weak base added is, for instance, ammonia, the conductometric titration with lithium hydroxide consists substantially in the neutralization not of H 3 0 + ions but of NH4+ ions. From this aspect the ammoniacal solution of an acid, HA, having as dissociamay tion constant in n-ater K , = be considered similar to the solution of an acid having as a dissociation constant,
+
1 2 3 LCLLMC. O F T I T R A R T
4
5
UL.
Figure 1. Comparison of conductometric titrations of vanillin with 2.48 1 N lithium hydroxide in absence and presence of weak bases Total solution volume, 1 2 0 ml. 1 . 9 5 % ethyl alcohol = 13 mi., sample = 1.059 gram, N = 0 2. Triethylamine = 6 ml., sample = 1.021 gram, N = 30 3. 1 4 . 7 6 N ammonia = 6 ml., sample = 0 . 9 4 5 gram, N = 3 0 Ordinate values for each titration curve are shifted b y an amount N which i s indicated for each compound
hoff’s method, for concentrations lower than 0.04 equivalent per liter (3, 4). I n this work, reference is made to the results obtained in performing conductometric titrations with lithium hydroxide, of phenols and other weak and very weak acids previously dissolved in aqueous solutions of ammonia or other weak nitrogen bases present in excess. The preliminary addition of an excess of a weak base has a double function. First, the presence of the weak base, B, shifts towards the right the dissocia-
+
neglecting the variation of activity coefficients due to the presence of an excess of ammonia. The considerable increase in the dissociation of the acid, H-4, not only leads to a marked improvement in the angle a t the equivalence point, so that frequently n here it was obtuse in the absence of a weak base i t now becomes acute, but also makes it possible to titrate groups not previously titratable by standard conductometric methods, as, for instance, salicylic acid and thiolignin. Secondly, because aqueous solutions of weak bases are generally good solvents for acidic compounds, water-insoluble compounds can usually be titrated without having to use mater-organic solvent mixtures n-hose adverse influence on the angle a t the equivalence point is knovx. During the titration, before the equivalence point, the hydroxyl ions of the titrant combine n-ith the B H + cations originating from the n-eak base, B. The variation of the conductance depends, therefore, on the difference between the mobilities of the Li+ and BH+ ions as \yell as on the variation of the concentration of the BH+ ion, which in turn deVOL. 31, NO. 6, JUNE 1 9 5 9
* 989
pends on the value of the dissociation constant of the weak base. After the equivalence point the conductance increases sharply as a consequence of the rapid increment of free OH- ions. The equilibrium
+ HzO C- NH, + OH-
NHs
existing with the equilibrium NHa
+ HA
NH:
+ A-
influences the curvature in the vicinity of the equivalence point without causing a shift of the point. I n fact, the titration considered here may be taken as the titration of a weak acid-e.g., phenolicin the presence of an extremely weak acid (water) whose neutralization line is constituted by the free base line. Because of the considerable difference between the two dissociation constants, there is no interference, as the results show. Among the weak bases used in these experiments (ammonia, pyridine, ethylamine, diethylamine, triethylamine) the best results were obtained with ammonia, because of the great mobility of the NH: ion. For comparison, Figure 1 shows three curves relating to the conductometric titration of vanillin. The first curve gives the results obtained using the standard method, while the second is with an excess of triethylamine, and the third with an excess of ammonia present. It appears from the graph that the conductometric titration in the presence of an excess of ammonia results in the titration of vanillin approaching the analytically more favorable case of the neutralization of a strong acid with a strong base. Thus the possibility of obtaining very good angles a t the equivalence point can be foreseen, by the application of the proposed variant, with acids even weaker than vanillin. I n fact, as seen from Figure 2, in the titration of three weak acids of phenolic or enolic type at the same sample concentration (0.055 equivalent per liter) the use of an ammonia concentration of 0.2 equivalent per liter is sufficient (when the compound solubility allows it) to obtain angles less than 90". Some small increases may appear, in some cases, after the minimum point; they can be attributed t o a variation in those characteristics of the solvent, which are a function of the ammonia concentration-that is, a variation of ionic mobilities. This is also demonstrated by the fact that the smallest initial resistance of the solution is observed in the experimental conditions corresponding to the minimum value of the enclosed angle. I t can also be seen from Figure 2, that the ammonia concentration values used in the titration can be varied within a wide range without greatIy increasing the size of the intersection angle. It has, 990
ANALYTICAL CHEMISTRY
125
'
)-
1
I
I
I
I I 2 loo
1-
i
I
i
I
m 75
w w
w
0
w 0
50
4
25
~~
~
1 2 3 4 A M M O N I A C O N C E N T R A T I O N EO. / L
5
Figure 2. Values for angles between neutralization line and free base line against ammonia concentration Sample concentration 0.055 equivalent per liter 1. Vanillln 2. Guaiacol 3. Acetoacetanilide
however, been observed that the straight portion of the neutralization line becomes shorter on increasing the ammonia concentration: it is therefore advisable, when possible, to maintain the ammonia concentration a t less than 2 or 3 equivalents per liter. In Figure 3, the values of the equivalence point angles are plotted against the sample concentrations; as could be predicted, the angles are considerably improved by increasing the sample concentration. I n this series of measurements, performed maintaining a constant ratio between ammonia and acid concentrations, the length of the first straight line remained practically constant. I n all the experiments, the conductometric titration gave quantitative results even a t the lowest concentration tried (0.012 equivalent per liter). EXPERIMENTAL
Procedure and Reagents. I n all titrations, 120 to 130 ml. of aqueous solution containing ammonia and the sample, in concentrations as indicated, were used. The determinations were carried out a t about room temperature. The ammonia was obtained, free from carbonate, by heating a reagent grade concentrated solution in the presence of barium hydroxide and collecting the gas in carbon dioxide-free distilled water. Reagent grade samples of acids were used. The monocarboxylic and monophenolic compounds were further purified by sublimation or crystallization until melting points corresponded strictly with those recorded in the litera+tutul.e. The titrations were carried out using an approximately 2.5N soIution of carbonate-free lithium hydroxide, stand-
I 1 I 0050 0075 010 SAMPLE CONCENTRATION EQ/L.
1
I
0025
0125
Figure 3. Values for angles between neutralization line and free base line against sample concentration Ratio between equivolentr of ommonla and sample = 15 1 . 2-Naphthol (23% ethyl alcohol solution) 2. Guaiacol 3. Phenol 4. Vanillin
,
1c3nr.
Figure 4.
Conductivity cell
ardized with potassium biphthalate in the presence of phenolphthalein. The reciprocals of the measured resistance values were corrected for the volume variations. Apparatus. I n carrying out the resistance measurements two different setups were used, the simpler of the two consisting of an alternating current Wheatstone bridge, the characteristics of which were: frequency 50 hertz; bridge supply 10 to 15 volts r.m.s., ratio of arms 1 t o 1; smallest graduation 1 ohm; sensitivity 1%. The bridge was not balanced for capacities, and the resistance balance point was obtained bv finding the smallest off-balance point, using amplifier, rectifier, and milliammeter. A normal immersion cell ( K = 0.68) was used with the bridge, and the solution was kept within 2~0.1"C. of the working temperature. When greater accuracy and resolving power were desired, a more elaborate setup was used. This was also used for the titration of compounds sensitive to air--e.g., hydroquinone, pyrogallol, etc.-because the operation could be carried out in an inert atmosphere of nitrogen. It consisted of an alternating
I
I
-___ _ _ ~
-1 Y ' -
-
~-
VOLUMEOF TITR~NT
Figure 6. Conductometric titration of polybasic acids in ammonia
Figure 5. Conductometric titration of monobasic acids in ammonia with 2.5N lithium hydroxide
Sample Concn., Ammonia Concn., Compound ( N ) MmoleslLiter Equivalents/Liter 1 . Resorcinol (- 30) 27.9 2.03 2. Hydroquinone ( 0 ) 63.5 0.28 3. Pyrogallol (50) 54.5 0.28 4. Alizarin ( 1 50) 27.8 1.54" 0.42 5. Phloroglucinol (- 80) 47.8 1.05 6. Salicylic acid ( 0 ) 50.8 7. o-Caumaric acid (100) 27.8 0.91 8. Catechol ( 1 30) 63.5 0.28 a Before adding aqueous ammonia it was necessary to dissolve the alizarin sample in a few milliliters of methanol Ordinate values for each titration curve are shifted b y an amount N, which i s indicated for each compound
Sample concentration 0.055 equivalent per liter Ammonia concentration approx. 0.9 equivalent per liter 1. Phenylacetic acid, N = 0 2. Benzoic acid, N = 30 3. Aminobenzoic acid, N = 63 4. 2-Naphthol, N = 0 5 . Guaiacol, N = 20 6. Phenol, N = 40 7. Vanillin, N = 60 Ordinate values for each titration curve are shifted b y an amount N which is indicated far each compound
current Kheatstone bridge constructed with parts marketed by the Pye and Trub-Tauber companies. The bridge was supplied through an insulation transformer. Balance point zero readings were made on a Trub-Tauber vibration galvanometer. Cell capacity was balanced by additional capacitors in parallel to the variable resistance. The characteristics of the bridge were: frequency 50 hertz; bridge supply 15 volts r.m.s. ; ratio of arms 1 to 1: smallest graduation 0.01 ohm; sensitivity 0.57,. The cell used with this bridge is schematically represented in Figure 4. The electrodes consisted of t n o small disks of platinized platinum (6 8 mm.) facing each other a t a distance of 10 mm. and held by glass supports which were part of the cell body. The solution was stirred by displacing and circulating the liquid between the two spherical vessels using a rubber bulb, P. A rubber bladder, V , prevented contact with air of the cell contents during stirring. The cell was immersed in a water bath thermostatically conC. trolled within the limits of =!~0.002~ of the working temperature. The microburet filled automatically and had a capacity of 5 ml. with 0.01nil. graduations. TITRATION OF MONOBASIC ACIDS
The application of the conductometric method as described to the titration of carefully purified carboxylic and phe-
___I
nolic acids has been satisfactory, the greatest errors being less than 1% of the theoretical value. The titration curves are recorded in Figure 5 . Enols and imides m-ere reagent grade and were titrated as received. Table I compares the analytical results obtained using the method described with those obtained using Fritz's determination ( 2 ) . The angles a t the equivalence point of the conductometric titration curves of the compounds listed in Table I are sharp and their values are very near to those obtained with phenols. The method vias inapplicable to the titration of ethyl cyanoacetate, because of an excessive opening of the angle which makes the location of the end point difficult and inaccurate. I n the titration of ethylmalonate, a t low ammonia concentrations, two intersection points are noted; the second of these corresponds to 100% of the theoTable
I.
Conductometric Titration
retical d u e . The first intersection point falls in different positions corresponding to 60 to 80% of the theoretical, according to the experimental conditions (ammonia concentration, time, temperature), and probably represents the neutralization of the enolic form which is present in the solution during the initial phase of the titration. TITRATION
OF
POLYBASIC ACIDS
I n some cases, a considerable excess of ammonia results in a shortening of the straight-line portion of the neutralization curve. This becomes much more noticeable n-hen the straight line relating to the neutralization of the second acid equivalent of a polybasic acid is considered. Because of this, very low concentrations of aninionia have been used in nearly all cases, except with compounds containing very weak acid functions-e.g., resorcinol.
of Enols and Imides in Ammonia
(Sample concentration = 0.055 equivalent per liter) % of Theoretical Value __ Ammonia Concn., Volumetric Conductometric Equivalent/ Compound titration (Fritz) titration Liter Acetoacetanilide 101.1 100.5 2.0 Methone (5,5-dimethyl-1,3cyclohexanedione) 100.4 99.7 1.4 1.4 Ethylacetodicarboxylate 98.4 99.1 Ethylmalonate 101.7 100.4 0.4 Succinimide 99.5 99.3 1.4 ~
~~
~~
VOL. 31, NO. 6, JUNE 1959
991
Tititration curves, relating to a group of eight phenolic compounds, are shorn in Figure 6. From these curves it TT ould appear that tTTo acidic groups have been distinguished in resorcinol, hydroquinone, pyrogallol, and, less sharply, in salicylic acid for which, to date, it has not been possible to reveal conductometrically any equivalence point ( 1 , 3). The values of titrant corresponding to the second equivalence point are, u-ithin experimental errors, double those corresponding to the first equivalence point. For phloroglucinol, alizarin, and o-coumaric acid (o-hydroxycinnamic acid), only one equivalence point was found, corresponding to the neutralization of try0 acidic groups. On the other hand, catechol reacted to give a single
equivalence point corresponding to one acidic group. The reproducibility of the results was verified by repeating the titrations several times. The application of this conductometric method to the titration of thiolignin samples revealed several acidic groups, information about which will be communicated later. CONCLUSION
The proposed titration method is suitable for a wider application than the original one. I n every case titration curves n i t h much more acute equivalence point angles and which allow a more accurate and sure interpretation of the results are obtained.
ACKNOWLEDGMENT
The authors are indebted to Angelo Terzi for the design of the thermostatic bath and the arrangement of the bridge. LITERATURE CITED
(1) Britton, H. T. S., “Conductometric Analysis,” TI-. G. Berl, ed., “Physical
Methods in Chemical Analysis,” Vol. 11. DD. 51-104 -4cademic Press, New Ydrk ,&1951. ( 2 ) Fritz, J. S., ASAL. CHEM.24, 674-5
(1952). (3) Kolthoff, I. M., IND.EKG.CHEM., d s a ~ED. . 2, 225-30 (1930). (4)Sarkanen, Ilyosti, Schuerch Conrad, ASAL. CHEX27, 1245-50 (1955).
RECEIVED for review September 15, 1958. Accepted December 22, 1958.
(Et hy lenedinit riIo)tet raacetic Acid Titration of MetaI Ions Polarized Mercury Electrodes A. E. MARTIN and C. N. REILLEY Chemistry Department, University of North Carolina, Chapel Hill, N. C. ,To extend existing methods for detection of end points in the chelometric titrations of metal ions, polarized mercury electrodes were used in the EDTA titrations of calcium(ll) and copper(l1). Anomalous reductions of mercury(l1)EDTA in ammonia buffers were investigated b y means of current-time curves. End point was detected in the titrations of millimolar concentrations of calcium(ll) and copper(l1) with EDTA. The advantages of polarized end point detection methods and suggestions for their use in continuous monitoring of metal ions are outlined. The methods may b e extended to include various metal ions using EDTA as well as other chelons.
P
methods of titration have distinct advantages, among which are rapidity and stability of electrode response, simplicity of physical components required, and often easily interpretable titration data. Consequently, the adaptation of polarized end point detectors for (ethylenedinitri1o)tetraacetic acid (EDTA) titrations of metal ions seemed highly desirable. Polarized systems, however, depend on the electroactivity of some species in solution and it is not immediately apparent how a polarized system of electrodes could be utilized for the chelometric titration of nonelectroactive metal ions such as calcium. OLARIZED
992
ANALYTICAL CHEMISTRY
The addition of another metal ion Fhich is electroactive and whose concentration bears a relationship to the titrate concentration makes possible this type of end point detection. T o be of general applicability the added component must have ideally, reversible electrode behavior, high stability of its chelates (even in acid solution), reduction potentials prior to oxygen interference, and no deposition of metals foreign to that of the electrode material. Coupling a mercury electrode with mercury(I1)EDTA yields a system which fills the above requirements admirably. The basic procedure is simple. To the metal ion to be determined are added a quantity of mercury-EDTA and a n appropriate buffer and the titration is carried out with standard EDTA solution. Before the titration begins, some of the mercury(I1) is displaced from its EDTA complex by the metal ion, the extent of replacement being a function of solution conditions (pH, buffer type and concentration, and concentration of the reacting species). A simplified equation representing the displacement is: Ll(X),
+ HgY + (y - p)X e Hg(X)v
+ MY
(1)
where h3: is the titrate metal ion, X is a complexing agent added in excess (nominally the buffer), and Y is EDTA (charges eliminated for simplicity). The species actually reacting a t the electrode is Hg(X), and it is the decrease in the
CEMENT
piI Figure 1. Electrodes for titration of metal ions
polarized
concentration of Hg(X), that is actually observed as the titration proceeds. Because Hg(X), concentration is directly related to titrate metal ion concentration, the titrate concentration is indirectly indicated. I n a somewhat similar way, Laitinen (5) employed zinc as an “indicator ion” to follow amperometrically the EDTA titration of calcium at high pH using the dropping mercury electrode. Because of the very negative half-wave potential of the zincate ion, complete removal of oxygen was necessary, a disadvantage not experienced when mercury is used. Fixed mercury electrodes of the type illustrated in Figure 1 can be polarized
