Conductometric Titration of Sulfuric and Hydrochloric Acids and Their


points of the two hydrogens of sulfuric acid have been definitely differentiated in aceticacid. The conducto- metric plots for this acid exhibit two b...
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Conductometric Titration of Sulfuric and Hydrochloric Acids and Their Mixtures in Anhydrous Acetic Acid TAKERU HlGUCHl and CARL R. REHM School o f Pharmacy, University o f Wisconsin, Madison, Wis.

lowing neutralization of the first hydrogen can be ascribed to the lower degree of dissociation of the alkali sulfate, which was progressively formed in the system, as compared t o the alkali bisulfate. Although straight-line approximations are used in the plot of extrapolation purposes, theoretically the actual curves are probably much more complex involving square root and higher power terms of concentration. I n practice, however, a linear plot seems to be a fair approximation in the concentration range studied. This explanation, based on the assumption that alkali sulfates are very slightly dissociated in acetic acid, is certainly a rea-

Sulfuric and h j drochloric acids and their mixtures have been titrated conductometrically in anhydrous acetic acid with several alkali acetates. The results represent the first time that the neutralization end points of the two hydrogens of sulfuric acid have been definitely differentiated in acetic acid. The conductometric plots for this acid exhibit two breaks corresponding to the neutralization end points of the first and second hydrogens. Differences in the plots obtained from titrations with different alkali acetates agree with predictions based on conclusions of earlier investigators. Conductometric titration plots of mixtures of sulfuric and hydrochloric acids in acetic acid with lithium acetate were found to exhibit three breaks corresponding to the neutralization of one of the two hydrogens of sulfuric acid, the hydrochloric acid, and the second hydrogen of sulfuric acid, in t h a t order. Quantitative estimations of sulfuric and hydrochloric acids in such mixtures, based on the three breaks, are in agreement with the amounts present.

I

X CONNECTIOS with other studies being carried out at these laboratories, conductonietric titrations of sulfuric and hydrochloric acids and their mixtures have been carried out in acetic acid. The results are of particular interest because they represent so far as known the first time that the neutralization end points of the two hydrogens of sulfuric acid have been definitely differentiated in this solvent. These findings, together with a rationalization of the ohserved facts, are presented a t this time not so much tiecause of possible direct analytical application of the conductometric method to the particular acids studied hut rather because they cast some light on the complex behavior of acid-base reactions in acetic acid and other solvents of lo^ dielectric constant which are enjoying wide usage in analytical fields. Apparently very little work has been done on conductometric titrations of acids in acetic acid. Perchloric and hydrobromic acid have been titrated conductometrically in acetic acid with sodium acetate by Kolthoff and Killman (4)and perchloric acid has heen titrated conductonietrically with a number of organic amines by Hall and Spengeman (3). -4lthough these investigators have shown the general feasi1)ilit.vof the method, there hap been very little sustained interest in this field.

1

+ +

>\IHSO4 HAC ~ I I ~ S OHAC ~

+

5

8

7

8

9

First break corresponds to conversion of sulfuric acid t o sodium bi-1 sulfate: the second break corresponds to conversion of sodium bi311lfate t o sodium sulfate

t 1000-

900

-+

4

Figure 1. Conductometric titration of 100 ml. of 0.0176.1.1 sulfuric acid in acetic acid with 0.500M sodium acetate a t 23" C.

The general type of conductometric titration plot yielded by sulfuric acid solution in acetic acid when titrated by an alkali acetate solution is shown in Figure 1. There are two significant breaks in the curve which correspond very closely to the stoichiometric reactions:

+ HzSO, + RIHSO4

3

ML. SODIUM ACETATE

SULFURIC ACID

MAC hIAc

2

M V.

-

(1) (2)

It is evident that the plot is different from the usual acid-base conductometric plots found in aqueous systems. The initial increase in the conductivity in the figure can be attributed to the higher degree of dissociation of the alkali bisulfate into current-carrying alkali ions and bisulfate ions as compared to the original sulfuric acid molecules. The drop in conductivity fol-

700

I

I

I

2

I

3

I

I

I

1

5

6

ML. SODIUM ACETATE

Figure 2. Potentiometric titration of 2 ml. of sulfuric acid (0.049M) with sodium acetate (0.0507M) in glacial acetic acid

408

V O L U M E 2 7 , NO. 3, M A R C H 1 9 5 5

409

boilable one. Because of its high central charge, it seems logical to expect, for example, the triple ion Sap+SOn-- to be less dissociated than the ion pair Na+HS04- The effective negative charge tending to bind sodium ions would appear to be considerably greater in the case of I than in the case of I1 because of the greater electrophilic nature of the proton as compared to the sodium ion. S04XaSOlH -

I

I1

Conductometric Determination of Sulfuric Acid in ..icetic Acid

Titrant NaAc LiAe

Ii ..\c (C5Hii)aNHAc

Potentiometric plots which are presently widely used in nonaqueous titrimetric determinations do not, in contrast to the conductometric plots, bring out this behavior. As shomm in Figure 2 the potentiometric titration curve of sulfuric acid in acetic acid with sodium acetate exhibits only a single break corresponding to the titration of the first hydrogen. From potentiometric response alone there is very little to indicate that the second hydrogen can be titrated. From a purely theoretical standpoint, the glass electrode, moreover, ip not necessarily the best indicator of the extent of these acid-base reactions. It can be shown that the extent of such reactions is roughly governed by the constant (3) 11here

Table I.

K , = dissociation constant of the acid Kb = dissociation constant of the base K H A=~ autoprotolytic constant of acetic acid Kob = dissociation constant of the resulting salt

Since the extent is not purely a function of the strengths of the acid and of the base alone, an acidity indicator is thus not necessarily the best indicator of these reactions. Further information as to the nature of this system can be obtained b y comparing the types of conductometric titration plots exhibited by different bases. The magnitude of the break in going from the bisulfate to the sulfate would be expected to be strongly influenced by the nature and size of the cations involved if the rationalization is valid. Thus, if potassium acetate were used to titrate the acid, a less pronounced break would be expected, since the larger pize of the potassium ion (as compared to the sodium ion) nould re-

Millimoles Millimoles 1st Sdded Hydrogen Found 1.72 1.74 0.889 0.904 0.587 0.586 0.586

0.580

0.586 0.580

0,580 0,580

0.492

0,497 0.583

0.58ij

Millimoles 2nd Hydrogen Found 1.73 0.903 0.588 0.586 0.520 0.510 .

.

I

...

duce the coloumbic forces iiivolved in the ion-pair forination. Similarly, the use of a base containing a very large orgaiiic ration might result in a curve in which the break is almost conipletely suppressed. With a tiase containing a relatively small cation, such as lithium acetate, one ~ o u l dexpect the break to be even m w e pronounced than that produced by sodium acetate. These predictions are based on the conclusions of Kolthoff and Willman (4). -111 of these expectations ivere boi~icout experimentally. In Figure 3, conductometric titration plots of sulfuric acid titrated (top t o bottom) with triamylaiiiiiioniuiii acetate, potassium acetate, sodium acetate, and lithium acetate are shown. Because of the limited solubilities of lithium and potassium sulfates in the system, it was found necessary to employ a lower concentration of the acid in obtaining these plots than that used in Figure 1. This resulted in somewhat poorer end-point breaks. It is apparent, nevertheless, that the degree of the first break derreased progressively in the order lithium > sodium > potassium > triamylamine. The point of break for the amine does not appear on the graph but consisted of a very slight positive break. I n Table 1; analytical result.- shown in Figure 3 are given for each end point. Although no seiious efforts were made to obtain highly accurate data, because of the type of extrapolation used as discussed previously, these results nevertheless indicate for most cases good stoichiometric relationships. 3IIXTURE O F SULFURIC 4YD HYDROCHLORIC ACIDS

Although it is impossible to distinguish the alkalimetric titration end points of miutures of hydrochloric and sulfuric acids by potentiometric means either in n ater or in acetic acid, acceptable end points can be detected bv use of conductometric means in the latter solvent. From the results of the present inveatigation, it appears that one of the t u o hjdrogens of sulfuric acid is first titi ated, then hydrochloiic acid, and finally the second hydrogen of sulfuric acid. The conductance curves obtained on titrating 0.11951 hydrochloric acid alone in acetic acid i\ith 0.200.V lithium acetate, sodium acetate, potassium acetate, and triam~~latiiinoiiiuii~ acetate are shown in Figure 4. Significant breaks &ere obtained in the case of trianiylammonium acetate and lithium acetate. A slight break was obtained with sodium acetate. These Iireakf correspond closely to the stoichiometric reaction

HCI I

1

1

1

/

,

,

1

1

1

I.0 2.0 HOLES ALKALI ACETATE/HOLES

I

1

I

I

I

,

30 ACID

Figure 3. Conductometric titrations of 100 ml. of 0.00586.M sulfuric acid in acetic acid with 0.200.M alkali acetates at 25" C. 9 Triamylammonium acetate

0 Potassium acetate +0 Sodium acetate Lithium acetate

The first break corresponds t o conversion of sulfuric acid t o the respective alkali bisulfates; the second break to the conversion of the alkali bisulfates t o t h e sulfates

+ SaAc

-+

SaCl

+ HAC

The iiiitial rise in conductance can be ascribed to the foiiii:ttion of the chloride salts xhich appeal to be more highly conducting than hydrochloric acid. The positive break obtained in the case of triamylammonium acetate indicates that this base is mole highly dissociated than the rhloride salt The absence of any discontinuity in the plot obtained with potassium acetate indicates that potassium acetate and potassium chloride are effectively dissociated t o about the same degree The negative breaks obtained in the case of sodium and lithium acetate indicate that these bases are probably less dissociated than their respective chloride salts in thefe systems.

ANALYTICAL CHEMISTRY

410

Determinations of hydrochloric acid made with the previously mentioned bases (except potassium acetate) are shown in Table 11. The slightly low results obtained were undoubtedly due to the high volatility of hydrochloric acid in acetic acid. Titration of several different molar ratios of mixtures of sulfuric acid and hydrochloric acid with lithium acetate are shown in Figure 5. The initial rise in conductance was due to the formation of bisulfate. The first break corresponds to the stoichiometric end point of removal of one of the two hydrogens of sulfuric acid. The subsequently greater rate of increase in conductance upon further addition of base indicates that the lithium chloride formed is somewhat more conducting (dissociated?) than lithium bisulfate. The second break in the conductance curve corresponds to the stoichiometric end point of the neutralization of the hydrochloric acid in the system. The decrease in conductance after the second break is explained in the same manner as before, the results of marked tendency for ion-pair formation between the sulfate ion and the lithium ions The amounts of hydrochloric acid and sulfuric acid found, calculated on this basis, agree reasonably n i t h the amounts added as shown in Table 111.

Table 11. Conductometric Determination of Hydrochloric Acid in Acetic Acid Titrant NaAc

~IillimolesAdded" 1.19

Millimoles Found 1.11 1 19 1.12 1 19 1 20 LiAc 1,lQ 1.10 (CaH1i)sNHAc 1.19 1.19 1.16 1.16 a Because of the high volatility of hydrogen chloride in acetic acid, it was difficult to maintain accurate concentration of the gas in solution.

Table 111. Conductometric Estimation of Mixtures of Hydrochloric and Sulfuric Acids in Acetic Acid with Lithium Acetate Sample 1

Millimoles of Acid Added 0.559 0.589 1.16 0.589 0.559 1.18

Acids Added HC1

2

HCl HzSO4 HCl HzSOd

3

Millimoles of Acid Found 1st H 2nd H 0.594 0.580 0:526 1.16 0.540 0:560 0.626 ... 1.12 1.20

GEKERAL DISCUSSIOS

These conductometric results are indicative of the vast difference between acid-base behavior in water and in nonaqueous solvents, especially of low dielectric constant. From purely theoretical considerations, titrimetric differentiation in water of sulfuric acid, bisulfate, and hydrochloric acid is impossible either potentiometrically or conductometrically, all three acids being effectively of the same strength. Yet, as is evident from the present study, relatively sharp end points can be detected for mixtures of these acids in acetic acid.

It

conductometric plots it was tacitly assumed that the conducting species were essentially simple ions, it must be recognized that, especially a t higher concentrations, complex ionic agglomerates play important roles in the over-all process of electrical conductc ance ( 2 ) . Because of the complexity of these particular systems, hoi? ever, their full and complete analysis is unfeasible. The simplified picture presented earlier appears to be sufficiently valid to provide a working hypothesis for analytical studies in these areas. EXPERIMENTAL

Apparatus. B cylindrical jacketed borosilicate glass vessel, 5 em. in diameter and 12 cm. deep, was used as the conductance cell. A fitted rubber plug was provided with holes for the electrodes, stirrer, and buret tip. The electrodes were brightly polished platinum plates with a cross-sectional area of 2.25 sq. em. Distance between the plates was less than 2 mm. The cell constant n a s determined using 0.01M potassium chloride and was found to be 0.06. I t was found necessary to thermostat the conductance cell because of an appreciable temperature

12 0

61

3.75t-

E 2.75 g250

2 MOLES A L K A L I ACETATE/MOLES

nci

Figure 4. Conductometric titrations of 100 ml. of 0.0119M hydrochloric acid in acetic acid with 0.2001M solutions of alkali acetates at 25' C. 3 Triamylanimonium acetate

+00

Potassium acetate Sodium acetate Lithium acetate

$225

"

200 175 1.80 1.25

IO0 I

It is of some interest t o note that potentiometric titrations of these same acids in acetic acid actually give only little indications of the reactions taking place, KO significant breaks are found which can be correlated to the individual end points. This behavior, as discussed before, is probably due to the fact that the extent of the neutralization reaction is not dependent only on the strength of the acid and the base, as in water, but also on the relative dissociative tendency of the salt which is formed. Although in the somewhat naive explanations of the observed

2

3

4

5

6

7 B 9 10 1 I ML. LITHIUM ACETATE

12

1 3 1 4 1 5

16

17

U

Figure 5. Conductometric titrations of mixtures of sulfuric and hydrochloric acids in acetic acid with lithium acetate at 25" C. Plots for titrations of mixtures of hydrochloric and sulfuric acids with 0 , 2 0 0 M lithium acetate are shown: I. 1.16 millimoles of hydrochloric acid and 0.589 millimole of sulfuric acid 11. 0.559 millimole of hydrochloric acid and 1.18 millimoles of sulfuric acid 111. ,0.5B9 millimole of hydrochloric acid and 0.589 millimole of sulfuric acid

V O L U M E 27, NO. 3, M A R C H 1 9 5 5 coefficient of electrical conductance in acetic acid. FVater from a constant temperature bath circulating through the cell jacket permitted temperature control to Zt0.2" C. of the desired temperature. A Leeds & Northrup conductance bridge (catalog No. 4866) was used to determine the conductance values. I n order to increase the sensitivity of the bridge a t high resistances a cathode ray oscilloscope (Type 304-H, Allen B. Dumont Laboratories) was connected in parallel with the null point galvanometer of the bridge. The null point of the bridge was determined from the resulting screen pattern of the oscilloscope. Chemicals and Reagents. ANHYDROUS ACETIC ACID. Reagent grade acetic acid was rendered anhydrous by refluxing over boron acetate ( 4 ) for 4 hours and subsequently distilling the anhydrous acid. The water content of acetic acid prepared by this method vias less than 0.027, by the Karl Fischer method. STASDARD SULFURIC ACID. A standard solution of sulfuric acid in anhydrous acetic acid was prepared by diluting absolute sulfuric acid (5)Tyith acetic acid. The molarity of the acid solution was determined by barium sulfate precipitation. STAXDARD LITHIUM,SODIUM, 4 N D POT.-ISSIC.\r ACETATE. ACcurately weighed quantities of these reagent grade salts previously dried overnight in a vacuum desiccator were dissolved in anhydrous acetic acid and diluted to volume. STASD.4RD TRIAVYL OYICM ACET-ITE.Commercial triamylamine was purified by distilling several times under reduced pressure, rejecting the first and last 20% of the distillates. Accurately Feighed quantities of the amine viere dissolved in anhydrous acetic acid and diluted to volume. The molarity of the resulting solution &-as checked by titrating with standard perchloric acid in acetic acid, using quinaldine red as the indicator. It was noted that acetic acid solutions of the triamylamine used developed a deep red color upon standing for several days. STASD-IRD HYDROCHLORIC ACID. Anhydrous hydrochlorir arid A as passed into cool anh>-drous acetic acid until fairly satu-

411 rated. The resulting solution was diluted with acetic acid and the molarity was determined by silver chloride precipitation. Frequent restandardization of this solution was found necessary owing to the high volatility of hydrochloric acid in acetic acid. Procedure. Accurately measured quantities of previously standardized acids were pipetted into the conductance cell and diluted to a volume of 100 ml. with anhydrous acetic acid. The solution was stirred until it had attained the equilibrium temperature of the thermostated vessel. Small increments of standardized base were added and the solution was stirred for about 30 seconds. The conductance reading was taken nrhen it became constant after stirring had stopped. ACKNOWLEDGMENT

This study was supported in part by the research committee of the Graduate School from funds supplied by the Wisconsin Alumni Research Foundation. REFERENCES

Eichelberger, W.C.. and La Mer, V. K., J . Am. Chem.

Soc., 55,

3633 (1933).

Fuoss, R. II.,and Kraus. C., Ibid., 55, 476, 1019, 2387 (1933). Hall, N. F., and Spengeman, TV. F., Trans. Wisconsin Acad. Sci., 30, 51-6 (1937). Kolthoff, I. XI., and Willman, A , , J . A m . Cheni. Soc.. 56, 1007 (1934).

Kunzler, J. E., ASAL.CHEM.,25,93-7 (1953). R E C E I V EJune D 9, 1954. Accepted November 18, 1954. Presented before the Division of Analytical Chemistry at the 126th Meeting of t h e . ~ M E R I C A X C H E M I C ASOCIETY, L h'ew York, September 1954.

Purification, Purity, and Freezing Points of Sixty-four American Petroleum Institute Standard and Research Hydrocarbons A N T O N 1. STREIFF, A U R A R. H U L M E , P H Y L L I S A. C O W I E , N E D C. K R O U S K O P , and FREDERICK D. R O S S l N l Carnegie lnstitute of Technology, Pittsburgh, f a .

The purification and determination of freezing point and purity are described for 64 hydrocarbons of the American Petroleum Institute Standard and Research series, including 11 paraffins, 3 alkyl cyclopropanes, 1 alkyl cyclopentane, 4 alkyl cyclohexanes, 24 monoolefins, 12 alkyl benzenes, 3 dicycloparaffins,3 dinuclear aromatics, 1 cycloparaffin-aromatic, and 2 olefin-cycloparaffins. Values of freezing points and cryoscopic constants are reported.

T

HE investigation reported is a continuation of the work of producing highly purified hydrocarbons of the API Standard and Research series ( 2 , 5-9). This paper describes the purification and determination of purity and freezing points of 61 hydrocarbons, which include 11 paraffins, 3 alkylcyclopropanes, 1 alkyl cyclopentane, 4 alkyl cyclohexanes, 24 mono-olefins, 12 alkyl benzenes, 3 dicycloparaffins, 3 dinuclear aromatics, 1 cycloparaffin-aromatic, and 2 olefin-cycloparaffin hydrocarbons. The final lots of material labeled -4PI Standard are sealed in vacuum in glass ampoules and made available as -4PI Standard samples of hydrocarbons, by the Carnegie Institute of Technology. (Twenty-seven of the Standard hydrocarbons are also available from the National Bureau of Standards, Washington 25, D.C.) T h e material labeled B P I Research is made available in appropriate small lots through the American Petroleum Institute Research Project 44 for loan to qualified investigators for the measurement of needed physical, thermodynamic, and spectral properties. Table I gives the names of the 64 compounds, the laboratories providing the starting material, details concerning the first and

succeeding distillations or other methods of purification, the character of the plot of the freezing point of the hydrocarbon part of the distillate as a function of its volume, and the volumes of the final lots of API Standard and Research material. The procedures followed in the process of purification and determination of purity were the same as those described in previous papers (3,5-9). Details of the distillation apparatus and operations also have been described ( 4 , IO). Figures 1, 2, and 3 show graphically the results of some typical distillations. Figures 1, 2, and 3 represent the cases where the purest material is, respectively, largely in the forepart of the distillation, in the middle of the distillation, and in the after part of the distillation. I n each figure plots are given for refractive index, boiling point, freezing point, and purity, as a function of the volume of the hydrocarbon part of the distillate. As emphasized in the previous reports, the blending of fractions of distillate for the preparation of material of the highest purity can be done safely only on the basis of the freezing points. Table I1 gives the following information for the compounds measured: the kind of time-temperature curves, whether freezing or melting, used to determine the freezing point; the freezing point of the actual sample; the calculated value of the freezing point for zero impurity; the value of the cryoscopic eonstant, determined from the lowering of the freezing point on the addition of a known amount of a suitable impurity ( 3 , 4 ) ;and the resulting calculated amount of impurity in the -4PI Standard and Research materials. ACKNOWLEDGMENT

Grateful acknowledgment is made to the organizations mentioned in Table I for their contributions of starting materials.