Page images
PDF
EPUB

CHAPTER VIII

STANDARD ACIDS AND ALKALIES

GENERAL PRACTICES FOLLOWED IN THEIR PREPARATION. SCHEMES OF MOLARITY, NORMALITY, AND FACTOR VALUES. PREPARATION OF STANDARD HYDROCHLORIC ACID AND STANDARD SODIUM HYDROXIDE SOLUTIONS

122. In the actual application of acidimetry-alkalimetry it is necessary to employ standard solutions of acids and alkalies, namely, those whose concentrations are definitely known. Now since it is often possible in the use of standard solutions of acids or alkalies to adapt the choice of concentration either to lessening the labor of standardization or to shortening the time of subsequent calculations, three practices are in vogue with respect to the preparation of standard acids and alkalies, the same schemes being also applied to standard solutions in general:

1. They are made up to exact fractional molarity (or normality), for instance 0.1 or 0.2 or 0.5 molar.

2. They are made up to an approximate fractional molarity, their exact values determined, and the solutions used as they are.

3. They are made up to an arbitrary factor value, usually selected so that burette readings of the standard acid or alkali will give directly the percentage of the unknown constituent being titrated.

123. With respect to the first practice of making up to exact molarity (or normality) it is necessary, in order to accomplish this, to make the solution up to a greater strength than is required, then determine its value, dilute accordingly to bring its value down to the exact figure desired, and finally determine its new value. As this procedure involves two standardizations with an intervening dilution, the second practice has come into vogue, i.e., of making the solutions up to approximate fractional molarity, determining their values exactly and then using them as they are.

In regard to the designation of solutions in terms of molarity or normality a few words are necessary. The scheme of molarity, which in point of time was the latter of the two to be introduced, is based on the use of the mole as the unit weight of reagent to be employed in defining standard solutions, namely, that weight of reagent which represents the molecular or formula weight in grams.1 The number of moles of reagent per liter of solution is the molarity of a solution. Thus one mole per liter of solution constitutes a molar solution; one-tenth of a mole per liter of solution a tenthmolar solution, etc. The capital letter M is used as an abbreviation for molar. The great advantage of this scheme is its freedom from all ambiguity and the ease which it contributes in the handling of analytical data, particularly data relating to oxidation-reduction reactions.

As examples of molar solutions, a molar solution of hydrochloric acid contains 36.47 g. HCl, of sulphuric acid 98.09 g. H2SO4, of oxalic acid 126.05 g. H2C2O4 · 2H2O, of potassium dichromate 294.20 g. K2Cr2O7, all per liter of solution.

124. The scheme of normality2 is based on the use of the gram equivalent weight as the unit weight of reagent to be employed in making up standard solutions, namely, that weight of reagent which corresponds to or brings into reaction one gram equivalent of an element, or more specifically that weight which corresponds to or brings into reaction 1.008 grams of hydrogen or 126.9 grams of iodine. The number of gram equivalents of reagent per liter of solution is the normality of a solution. Thus one gram equivalent of reagent per liter of solution constitutes a normal solution; one-tenth of a gram equivalent per liter of solution a tenth-normal solution, etc. The capital letter N is used as an abbreviation for normal. The idea back of this scheme is that if solutions are made up to contain the same number of

1 The use of term mole was proposed and introduced by Ostwald. It is of very great importance in that it brings out very clearly many fundamental chemical relationships which otherwise would be scarcely discernible. In this regard see any textbook on Physical Chemistry.

2 According to A. Classen, Friedrich Mohr's Lehrbuch der Chemisch-Analytischen Titrirmethode, Fr. Vieweg u. Sohn, Braunschweig, 1896, 7th ed., p. 56.

So far as is known this scheme was first introduced by John Joseph Griffin of London about 1860 in the effort to bring about a uniformity in the employment of standard solutions which before that time had been used without reference to any special system. It was adopted by Mohr and largely extended by Mohr during the period.

gram-equivalents of reagents per equal volumes, then for many of these solutions, though not for all, 1 c.c. of the one solution will furnish just the amount of reagent demanded by 1 c.c. of the other. The great disadvantage of this scheme is the ambiguity to which it often leads because of the fact that the weight of reagent necessary to make a normal solution depends upon the purpose for which it is used, thus in the case of phosphoric acid a normal solution may contain either 32.71 g., 49.05 g. or 98.09 g. H3PO4 per liter of solution while a normal solution of potassium dichromate may contain either 73.55 g. or 48.03 g. K2Cr2O7 per liter of solution, consequently in these cases it would be open to doubt as to what is meant by a normal solution unless some explanatory note is given in the text. The scheme of normality as a basis of definition is inadequate and should be abandoned. For hydrochloric acid (HCl) a normal solution contains 36.47 g. HCl per liter of solution, because this weight corresponds to 1.008 g. of hydrogen; for sulphuric acid (H2SO4) a normal solution contains one-half of 98.09 g. or 49.05 g. H2SO4 per liter of solution; for phosphoric acid (H3PO4) it is a matter open to doubt. 125. With respect to the practice of making up standard solutions to an arbitrary factor value, this is particularly suited for routine work when the factor value is so selected that the burette readings of the standard solution will give directly the percentage of the unknown constituent being titrated. We will now show how the selection of such a factor value is arrived at.

Let p f

n

=

=

% constituent to be represented by 1 c.c. of standard solution
value of the standard solution in terms of the constituent being
determined, i.e., the number of grams of constituent corresponding
to 1 c.c. of the standard solution

= number of c.c. of standard solution used in the titration

[blocks in formation]

F. P. Treadwell, Analytical Chemistry, translation by W. T. Hall, Vol. II. Quantitative Analysis, John Wiley and Sons, New York, 1913, p. 532, calls attention to this fact.

G. Lunge, Technical Methods of Chemical Analysis, translation by C. A. Keane, Vol. I. Part I, p. 81, Gurneyand Jackson, London, claims that in figuring the weight of reagent cor

It is evident from the foregoing relationship that if we arbitrarily fix the value of p, we can vary ƒ and W; if, however, p and W are fixed, there is only one value of ƒ that will satisfy the equation.

Example. Suppose that it is required to make up a standard hydrochloric acid solution of such a value that 1 c.c. of it shall correspond to 2% sodium carbonate, when 1 gram of the sample containing the sodium carbonate is taken for analysis, and methyl orange is used as the indicator in the titration. We have

[blocks in formation]

namely, 1 c.c. of the hydrochloric acid solution shall be equivalent to 0.0200 g. Na2CO3, or from the reaction:

Na2CO3 + 2HC1→ 2NaCl + H2O + CO2
106.00 2(36.47)

1 c.c. of the standard solution shall contain 0.01376 g. HCl.
The plan of adjusting the weight of the sample to the strength
of the solution is particularly advantageous in routine work where
use is made of a standard solution which changes in strength and
has to be frequently restandardized. In any case, however,
where this system is to be used, it is necessary that the sample
to be weighed out is in a finely pulverized condition, otherwise
with a coarsely divided substance, like iron drillings for instance,
one might easily lose more time in adjusting the exact weight
of material on the balance pan than would be saved in the cal-
culation.

126. Another scheme which is often used in routine work to express the value of a standard solution is in terms of the constituent which is to be determined. This scheme is convenient in that the weight of the constituent is found at once by multiplying the number of c.c. of standard solution used by the value of 1 c.c.

responding to 1.008 g. of hydrogen it should be figured in terms of the substance to be examined and says that it is only accidental when the ratio of the gram equivalents refers to the composition of the normal solution itself. On this basis a normal solution of phosphoric acid would contain 98.14 g. H3PO4 if figured in terms of the sodium hydroxide required for titration with methyl orange as indicator, or 49.07 g. H3PO4 if phenolphthalein were used as indicator.

By the value of a solution we shall mean any measure which either directly or indirectly gives the weight of reagent which the solution contains per unit volume.

127. The scheme of expressing the value of a standard solution in terms of the dissolved reagent is not much used because a great part of volumetric analysis consists in the routine determination of some given constituent which is different from the dissolved reagent, and therefore this would necessitate for each titration, in order to arrive at the weight of the constituent sought, an intermediate calculation of the weight of standard reagent corresponding to the volume of standard solution used, a roundabout procedure which can most always be abridged by employing some one of the other schemes.

[ocr errors]

In conclusion, as to the various ways of expressing the values of standard solutions we would remark that while it is a matter of choice as to which scheme shall be employed, it is usual to select that one which will lend itself most readily to a simplification of the subsequent calculations.

128. Standard Acids. There are practically only two acids which are suitable for general use as acidimetric standards, namely, hydrochloric and sulphuric. Nitric acid, although it is likewise a strong acid, is not desirable because it almost invariably contains a small percentage of nitrous acid which has a destructive action on the two important indicators, methyl orange and methyl red.

The primary standard that is used volumetrically to standardize either hydrochloric or sulphuric acid is sodium carbonate in conjunction with either methyl orange or methyl red as indicator. The sodium carbonate must be prepared as described herewith.

129. Preparation of the Sodium Carbonate. - Place about 5 grams of pure sodium bicarbonate in a platinum or a glazed porcelain crucible, and put the crucible with its contents on a sand bath, piling the sand about halfway up the sides of the

5 Hydrochloric acid is also standardized gravimetrically by the following methods: (a) by precipitation as silver chloride, see § 172; (b) by neutralization with ammonium hydroxide, evaporation to dryness of the ammonium chloride and determination of the ammonium chloride by loss of weight upon volatilization. The first method although the method used for atomic weight determinations does not give good results in the hands of students, the results being from 5 to 20 parts per 1000 low. The second method requires the use of large platinum dishes (of about 80 to 100 c.c. capacity) so that its general use is largely restricted.

6 Sulphuric acid is also standardized gravimetrically by precipitation as barium sulphate. 7 Instead of a sand bath, an air bath may be used. A most excellent and simple air bath is made by taking a silica beaker about 6.5 cm. diameter by 8 cm. high, fitting it about 1 cm. from its top with a wire triangle made of some difficultly oxidizable wire (preferably manganin), so arranging the construction that the sides of the triangle extend a suitable length to engage a 9 cm. iron ring by means of which the beaker is suspended from a ring stand. Inside of the

« PreviousContinue »