An Experiment in Acidimetry JAMES M. HENDEL Hunter College, New York, New York
IN
TEE titration of an acid by standard alkali the graph showing how the pH of the acid splution changes with the addition of base may be used to locate the end point, which is taken as the mid-point of the portion of the curve parallel to the pH axis, or the point where the slope of the curve is maximum. The purpose of making such a graph is, of course, to guide one in the selection of the proper indicator for the particular acid titration. The method of determining the pH is usually electrometric, originally the hydrogen eiectrode, or latterly the glass electrode, being used in conjunction with a standard calomel electrode. The graphical results of such experiments may be obtained by a much simpler method requiring no experience in electrical measurements, which therefore could be introduced early in the course in quantitative analysis. This method involves the use of a mixed indicator solution, sometimes called a "universal" indicator. The curves obtained by the two methods are similar in form and practically coincide in the critical range between pH 5 to pH 9. The indicator method proves to be just as valid as the electrometric in judging which individual indicators are suitable for the titration in question. The universal indicator solution1 used in these experiments consisted of: Methyl orange. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 0.05g. Methyl red.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 0.15 g. Bromthymol blue. . . . . . . . . . . . . . . . . . . . . . . . . . . 0.30 g. Phenolphthalein.. . . . . . . . . . . . . . . . . . . . . . . . . . . . 0.35 g. 66% Ethyl alcohol. . . . . . . . . . . . . . . . . . . . . . . . . . 1 liter
of the next unit pH value and the volume of added base plotted against this value. Successive portions of base were then added to produce the colors shown by the buffersof known pH value, proceeding cautiously between pH 5 and pH 9 where a drop or two of base often was sufficient to raise the pH by one whole unit. The buffer solutions used to produce the colors for comparison with the unknown solution were mixtures of 0.2M NalHPOl with 0.1 M citric acid as follows: PH 9 0 4 5
76 8
Phosphate, a on ml. 0. AL
15.42 20.60 25.26 32.94 38.90
-.
Citric w ma i d , ml. OI.10
24.58 19.40 14.74 7.06 1.10
To each of these was added 0.40 ml. of universal indicator. To prepare the buffer for pH 9, 50 ml. was taken of a solution which was 0.2 M with respect both to HaBOa and KC1, to which was added 21.40 ml. of 0.2 M NaOH, the mixture being then diluted to a volume of 200 ml.; for pH 10, 50 ml. of the boric acidpotassium chloride mixture was added to 43.90 ml. of 0.2 M NaOH, and diluted to 200 ml. To these two solutions, 2 ml. of universal indicator was added. Preliminary experiments showed that the titration curves for hydrochloric and for acetic acid resembled in form the typical curves for a strong acid and weak acid as determined by electrometric methods. To show their identity the course of the colorimetric titration was followed with glass and saturated calomel electrodes. The tables and graphs presented here The colors exhibited by this indicator are: showthe results obtained with the two indicator systems R e d . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . at pH 3 or less for acetic, succinic, and acids, where only one Red orange.. . . . . . . . . . . . . . . . . . . . . . . . pH 4 Orange . . . . . . . . . . . . . . . . . . . . . . . . . . . . . pH 5 inflection point is obtained, and for phosphoric acid Yellow.. . . . . . . . . . . . . . . . . . . . . . . . . . . . pH 6 where two inflection points are to be expected. The Yellow-green.. . . . . . . . . . . . . . . . . . . . . . . pH 7 reverse titration, of sodium carbonate by standard Green-blue.. . . . . . . . . . . . . . . . . . . . . . . . pH 8 HC1, likewise showed close coincidence in the values Blue . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . pH 9 by the two methods. Violet.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . pH 10 Values in Table 1 from a pH of 5 to 11 as found by the Red-violet . . . . . . . . . . . . . . . . . . . . . . . . . . pH 11 or more indicator method, are plotted in Figure 1. This plainly The procedure in generalwas to take 25 ml. ofa~proxi- shows an inflection point a t a pH of 7 and the practical mately 0.1 N a c 4 add 1 ml. of universal indicator identity of the values by the two methods, since the and 75 ml. of distilled water, and compare the color with ele~trometriomeasurements as squares on the that of a buffer solution of known pH containing the graph were reliable only to one-tenth of a pH unit. ~ h , important values in ~ ~2 are b l ~in same proportion of indicator. Sufficient 0.1 N sodium hydroxide was then added to bring the color up to that lqgure2. ~h~ maximum slope of this curve ml. I srnTH, T. B., 'cAna]ytioal prooesses,~~~ d ~ ~ ~ & ~ ~a t 24.85 ~~ dl and d a pH of 9. A titration using phenolGo., London. 1929, p. 363. phthalein required 24.90 ml. of base. But any other 148
MARCH, 1952 TABLE 1 Titration of 25 ml. of 0.0832 NAcetic Acid by 0.1004 NNaOH
M1 of NaOH
Colorimet+
0
pH
EleetromlricpH
3
3.3
Fi2. Succinio Acid Titration Curve (Squkrea indi~atedeotromatric values)
There are evidently two maxima in the curve, around pH 5 where 6.65 ml. of base were used for neutralization of the first hydrogen ion and near pH 10 corresponding to 13.45 ml. for the second hydrogen ion. A separate titration with phenolphthalein required 13.33 ml. of base. A CLASS EXPERIMENT , L 19
20 i
1
MI.
21
22
NaOH
Aseti. Asid Titration C u m
With a view to introducing a little of the "research" attitude into the course in quantitative analysis and to illustrate the value of titration curves in deciding what indicator should he used for a particular acid-base
(Squares indicate eleotrometrio valuea)
indicator changing within the range of a pH of 8 to 9.5 should prove useful. In Table 3 the identity of the pH values by the two methods is obvious. The titration should require an indicator changing between pH 8 and 9. Two titra.05 ml. of tions using phenolphthalein used 30.85 base showing the neutralization of all three hydrogen ions. As a more rigorous test of the validity of the indicator method, titrations of phosphoric acid were carried out to see if the two inflection points corresponding to the neutralization of the first two hydrogen ions would appear. Table 4 and Figure 4 show the results.
*
TABLE 3 Titration of 10 ml. of 0.300 N Citric Acid by 0.0973 NNaOH Ml. o j NaOH
Colorimetric pH
Eleclrometrie pH
0 12.30 19.30
3
2.5 4.0
4 5
5.3
TABLE 2 Titration of 25 ml. of 0.1000 N Succinic Acid by 0.1004 N NaOH Ml of NaOH
Colwimetrie pH
0 4.20 15.10
4
3 5
Electrometric pH 3.3 3.9 5.3
MI. NaOH Figun 3. Citric Aeid Titration C u m (Squsres indicate electrometrio values)
150
JOURNAL OF CHEMICAL EDUCATION
of
z5 ,,I.
MI.of NaOH 0 4.00 6.00 6.45 6.65 6.80 7.15 10.00 12.35 12.55 13.30 13.40 13.55 15.00
TABLE 4 H
of 0.0259
Colorirnetvie pH 3
... 4
... 5 ... 6 7
. .. 8 9
...
10 11
10
1 ml. of universal indicator was added, and the titration ~ by ~ 0.0913 o ~ M N ~ O H was carried out, adding sufficient base each time to Elect~ornetriepH 2.8 3.2 3.8 4.1 5.1 5.6 6.0 7.0 7.8 8.0 8.9 9.5 9.6 11.0
M XIMUM
s
OPE
MI. NaOH
Fi-
4. Phosphoric Acid Titration C u m (Squares indicate electrometric vsluea)
titration, each student was given a different acid and told to make 250 ml. of solution approximately 0.1 N. Since some of the acids were not soluble to that extent in cold water, these were dissolved in 95 per cent ethyl alcohol or in hot water to make a concentration requiring about 25 ml. of standard 0.1 N NaOH. In the usual case 25 ml. of the 0.1 N acid was diluted to 100 ml.,
change the pH by one unit as shown by the indicator. The comparison buffer solutions were provided by the instructor, made up in "wholesale" quantities. From the resulting titration curve the student was required to select three indicators which might be used to locate the end point, and to determine, by titrating, which one of these was the best. The indicators supplied were bromphenol blue, methyl red, bromcresol purple, bromthymol blue, phenol red, cresol red, thymolphthalein, and thymol blue. The use of methyl orange or phenolphthalein was discouraged since one of the minor objectives of the experiment was to de-emphasize these two indicators. Curves showing unmistakable inflection points were obtained for the following acids:
The curves for phthalic acid, Kl = 1.3 X lo-=, and tartaric acid, K I = 1.1 X showed a vertical portion in the curve only after neutralization of the second hydrogen ion, pH 8-9 for the former and pH 7-9 for the latter. Boric acid, K, = 6.4 X 10-10, with 25 ml. of glycerol added to 25 ml. of 0.1000 M solution showed a maximum slope between pH 8-9 corresponding to 24.06 ml. and 24.60 ml. of 0.1043 N NaOH; the theoretical volume for the first hydrogen is 23.97 ml. of base. A titration with phenol red (7.07.4) required 23.83 ml., with cresol red (7.6-8.2) 24.02 ml. and with phenolphthalein 24.03 ml. of base. For comparison with the experimental curve it is of value to plot the theoretical graph. For weak monobasic acids the ionization constant will yield the initial pH, the hydrolysis formula,for the salt will give the pH a t the stoichiometric point, and for intermediate points the equation pH = pK.
+ log
(=It) (acid)
will serve the purpose. This correlation of experiment and theory arouses considerable interest on the part of the student; it is not just another recipe.
