February, 1925
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Standardizing Acids and Bases with Borax’ ,
By M. G. Mellon and V. N. Morris PURDUSUNIVERSITY, LAFAYETTE, IND.
T
HE concentration of standard acids is most often de-
made in order to establish the actual value of borax as a standtermined by a direct comparison of the solution with ard. sodium carbonate as the primary standard. ConAs pointed out previously,’ electrometric titration curves sidering the requirements2 which such a substance must meet furnish valuable information regarding the course of the in order to serve both as a satisfactory primary reference reaction involved, and aid one in selecting a suitable indicator standard and as a regular working standard in titrametric for titrations. One curve similar to those reported here has analvsis. Dodge3 has tabulated the been determined8-that for borax and hydrochloric acid-but there defects of varyous compounds, inElectrometric titration curves show that is little information regarding the cluding ‘Odium which good end points may be obtained when using details of the work and no discushave been proposed for the standborax as a means of standardizing directly sion of the significance of the results. ardization of solutions for acidicertain strong acids. When used i n con-
metric and for deterjunction with sufficient mannitol, borax Reaction of Borax with Acids and minations. may also be used for the direct standardizaBases Having in mind the deficiency tion of strong bases. For titrating acids the Being a salt of a strong base and of sodium carbonate as a standard, pH range of the indicator used should be Lindfors4 has recently called atvery weak acid, borax shows an approximately that of methyl orange, and alkaline reaction in an aqueous tention again to the possibility of using borax, NazR40,.10Hz0, as for bases that Of phenolsolution, the extent of the hvdrolphthalein. ysis being sufficient for phenola basis of reference in checkinn phthalein to show a pink color in the solutions of acids. He states thac although borax is a hydrated compound, “in the hands of solution. When using borax as a means of standardizing men with little analytical experience or scientific education, it hydrochloric acid, for example, the reaction may be repreis probably the most accurate check available.” About the sented by the equation : same time von Bruchhausen5 recommended borax both as a NazB407 2HC1 5H20 4HsBO3 f 2NaCI primary standard for checking acids and as a suitable mateAs the neutralization proceeds the titration of a strong acid rial for ordinary use as a standard basic solution. He states that the salt is perfectly stable when properly prepared. in the presence of a weaker one is involved, as noted in an Dodge3 lists as the only objectionable feature connected earlier paper.9 One-half of the molecular weight of the [sowith its use for such purposes the possibility of the degree dium borate decahydrate is its equivalent weight. When using borax for standardizing a base, one has a of hydration being uncertain. Blasdale6 states that it is readily obtained pure by recrystallization and gives accurate choice of two procedures. The first and shorter method is an application of the procedure recommended by Tanana results in the standardization of acids. In one respect borax is one of the most interesting materials and Zuckermann’o for the determination of alkalies either available for standardizing solutions to be used in neutraliza- alone or in the presence of carbonates. As previously stated, tion reactions. In the references already noted the compound an aqueous solution of borax reacts alkaline. On the addiwas suggested for use in determining the concentration of tion of certain polyhydroxy organic compounds to the soluacids directly. The authors are unaware that any one has tion, such as glycerol or mannitol, its pH value rapidly drops previously called attention to the fact that it may also serve below 7, owing to the formation of a complex acid of much for the direct standardization of bases. For checking bases stronger acidity than boric acid.ll This solution may then we are not so limited in our choice of standards as in the be titrated with a base, using phenolphthalein as the indicator. case of acids, several excellent materials, such as potassium In this case the results obtained indicate that one-half the acid phthalate, being available. But in borax we have a single molecular weight of the borax is its equivalent weight. The second method involves a preliminary neutralization compound which can be used for the direct standardization of the borax by means of a strong acid. It is evident from the of both acids and bases. The object of the present work is to call attention to this foregoing equation that an amount of boric acid equivalent availability of borax as a primary standard for both acids and to the borax is liberated. By adding glycerol or mannitol bases, and to show the electrometric titration curves obtained to the solution this acid may then be titrated with a strong when using it as a means of standardizing various solutions. base according to the usual method for determining boric This report should be considered as only a preliminary con- acid. In the case of glycerol the additional reactions may be tribution, since no study was undertaken with the aim of represented by the equations :12 comparing borax with other standards, such as sodium carHaB03 f CaHs(0H)s ---+ CaHa(OH).HBOs 2Hz0 bonate. The work of other investigators upon this point is C3H6(0H).HB03 NaOH CsHa(OH).NaBOaf Hz0 accepted for the present. A careful comparison should be It is apparent in this case that the equivalent weight is one-
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1 Presented under the title “An Electrometric Study of the Use of Borax as an Acidimetric and Alkalimetric Primary Standard” before the Division of Industrial and Engineering Chemistry a t the 68th Meeting of the American Chemical Society, Ithaca, N. Y.. $eptember 8 to 13, 1924. * McBride, J . A m . Chem. Soc., 34, 393 (1912). 8 THISJOURNAL, 7, 29 (1915). 4 I b i d . , 15, 1046 (1923). 6 Arch. Pharm.,261, 22 (1923). 4 “Quantitative Chemical Analysis,” 1914, p. 294.
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Mellon and Morris, THISJOURNAL, 16, 123 (1Y24). Van Liempt, Rec. &as. chim., 39, 358 (1920); Z. anoi‘g. Chem., 111, 151 (1920); Hildebrand, J . A m . Chem. 5’06.. 35,847 (1913). 9 Mellon and Morris, Proc. Indiana Acad. Sci., 33,85 (1923). 10 J . Russ. Phys.-Chem. SOC., 41, 1469 (19091. 11 Abegg, “Handbuch der anorganischen Chemie,” Vol 11, Pt. 1, 1908, p. 311. 1 9 Mahin, “Quantitative Analysis,” 1924, p. 204. 7
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INDUSTRIAL A N D ENGINEERING CHEMISTRY
146
fourth the molecular weight of the borax, since four molecules of sodium hydroxide are required to react with the boric acid liberated from one molecule of borax. Present Work
The apparatus and the general method of procedure used in determining the p H curves were essentially the same as those described previously.' Titration curves were determined for
Vol. 17, N o . 2
what less highly ionized, could hardly be standardized with satisfactory accuracy. It should be noted that the method of standardizing this acid, with methyl orange as the indicator, made the solution fifth normal with respect to the first hydrogen. The relatively weak, monobasic acetic acid shows a curve still less satisfactory. Borax shows a good curve when titrated with sodium, potassium, or barium hydroxide in the presence of mannitol. A similar curve, shown in an earlier 14
f2 //
3
2
4
8
/2 /C 20 44 26 32 m / of tiPatin9 s o / u f i o n
3L
40
4/1
90
52
Figure I-pH Curves Showing Course of Neutralization for 30 MI. of Various 0.2 N Solutions Nos. 1, 2, 3, and 4 are for hydrochloric, sulfuric, phosphoric, and acetic acids, respectively, when titrated with borax. Nos. 6, 6, and 7 are for borax when titrated in the presence of mannitol with barium, potassium, and sodium hydroxides, respectively. Borax solution contained 7 to 8 grams of mannitol.
hydrochloric, sulfuric, phosphoric, and acetic acids when reacting with a solution of borax; for sodium, potassium, and barium hydroxides when reacting with a solution of borax containing mannitol; and for solutions of borax containing different amounts of glycerol or mannitol when reacting with sodium hydroxide. Borax was recrystallized416by allowing a hot, practically saturated solution to drop through a filter paper into a dish placed in cold water, with constant stirring. The crystals were separated on a Buchner funnel and air-dried on filter paper. Constant boiling point hydrochloric acid was weighed directly. The sulfuric and phosphoric acids were standardized against sodium carbonate, with methyl orange as the indicator, and sodium hydroxide free from carbon dioxide was used for the acetic acid, with phenolphthalein as the indicator. The barium and potassium hydroxides were also free from carbon dioxide. Fifth normal solutions were used throughout. The results are shown in the figures. Discussion
An examination of Figure 1will reveal the type of curve obtained for the various substances titrated. It is well known that under ordinary conditions only those substances can be titrated accurately whose p H curves for the titration have a vertical portion of sufficient length. It is evident that hydrochloric and sulfuric acids, representing, respectively, highly ionized mono- and dibasic acids, g!ve the proper type of curve with borax. Tribasic phosphoric acid, some-
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8
/2
/C
20
24
28
32
36
40
44
48
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m/ of fitroi'ing solufion Figure 2H Curves Showing Effect of Different Concentrations of Glycerol a n c f w a n n i t o l when Titrating a 0.2 N Solution of Borax with a 0.2 N Solution of Sodium Hydroxide
paper,' is obtained if boric acid is titrated in the presence of mannitol after its liberation from borax by means of hydrochloric acid. Electrometric titration curves enable one not only to ascertain whether a given neutralization reaction will give a sharp end point, but also to select an indicator with the proper range for the titration. I n general, the effective range of a suitable indicator should be opposite the steep portion of the curve. An inspection of the various curves considered as satisfactory will indicate that methyl orange or bromophenol blue, with effective ranges extending from pH 3.0 to 4.6, should be a suitable indicator for either hydrochloric or su1furic acid. Likewise, phenolphthalein, with a pH range of 8.3 to 10.0, is satisfactory for locating the proper end point when using borax as a means of standardizing bases. Assuming that the degree of hydration of borax can be controlled under suitabIe conditions, an objection to its use in certain cases may still be found in the fact that the liberated boric acid cannot be removed by boiling, as in the case of carbon dioxide liberated from sodium carbonate. The advantages of using borax as a standard, as compared with sodium carbonate, are its higher equivalent weight and the possibility of its being used for both acids and bases.
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