182
INDUSTRIAL A N D E.\-GIh’EERISG
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CHh’JIIISTRY
Vol. 15, KO. 2
tion Curves for Some C Acids and Bases as Determined by the Hydrogen Electrode’ By Clarke E. Davis, Earle T. Oakes, and Henry M. Salisbury NATIONAL BISCUITCo., 85 NINTHAVLE , New YORK, N. Y.
HYDROGEN ELECTRODESINCE the work of HilI t has been suggested that a complete set of electrometric titration The authors used a platidebrand2 and others, curves for the various acids and bases be made available through the num rod coated with platithe methods, theory, adoption of a standard method of procedure. A set of electrometric num black. This gives reliand apparatus for hydroable service over a long titration curves in which the same standard procedure was employed period of time if it is caregen-electrode t i t r a t i o n s throughout has been given for hydrochloric acid, sulfuric acid, o-phosfully handled. The platinum have been advanced to a phoric acid, primary calcium phosphate, and citric acid, with sodium electrode required from 3 to c o n si d e r a b 1e e x t e n t. hydroxide and sodium carbonate. Owing to the differencesin chemical 5 min. between additions of Coupled with this progress the titrating reagent to come reaction in some cases as well as differencesin the salt content of the t o equilibrium. The elechas been a gradually insolution, reversing the order of titration docs not result in titration trodes were cleaned after creasing demand on the curves that are mirror images of each other. each titration by immersing part of theoretical investiin sulfuric acid-dichromate gators, instructors in chemsolution and washing in distilled water, after which they were boiled three times in fresh istry, analysts, and practical technicians, for complete titration curves of each of the different acids and bases. In portions of distilled water. STIRRING-Efficient stirring is absolutely necessary to attain order for such curves to be of the maximum value all equilibrium quickly. The authors encountered difficulty in titrations should be carried out under the same conditions using a mechanical stirrer operated by an electric motor. of volume, temperature, etc. Furthermore, the procedure Stray currents caused such a variation in the observed voltages should be such that it conforms to correct analytical prac- of the cell that the mechanical stirrer had to be abandoned in tice, the results of which will give the analyst a complete favor of a hand-operated one. CALOMELELECTRODES-The saturated potassium chloridedescription of every stage of the operations he so often percalomel electrodes and the saturated potassium chloride salt forms. It will aid him in selecting the proper indicators and intelligently judging end-points. During the course of some researches in this laboratory it 6QDA LIME TUlE became necessary to titrate several acids and bases with the hydrogen electrode. Realizing that a great many other investigators must be doing the same thing for other acids and bases, we are publishing our methods and results in detail, with the hope that some standard may be adopted wherever possible, and that ultimately a complete set of titration curvw for all acids and bases will be available in such shape that they may be compared directly one with I I I li COTTON the other.
S
B$
PROCEDURE The procedure that we have adopted and that is proposed for general adoption is as follows: VOLUME OF SOLUTIONS Us&-Make up the volume of the solution to be titrated to 200 cc. and titrate with a reagent of such strength that 30 to 40 cc. are required for complete neutralization. Use a 50-cc. buret. The remaining 10 to 20 cc. are used t o finish the curve. This procedure is general quantitative analytical practice. TYPEOF CE&L-’I’he titration cell should be of the closed type. This is absolutely necessary; otherwise, owing to the length of time required for titrations, the various changes in the composition of the air of the laboratory will affect the solution undergoing titration and the result will be worthless. The authors found it impossible to duplicate titration curves in an open cell. Various types of titration cells have been described in the literature and many of them have been made available by the instrument makers. An exceedingly simple and efficient cell, used by the authors, is shown in Fig. 1. The solution to be titrated is contained in a 300-cc. flask or beaker, A close-fitting, 4-hole stopper carries the salt bridge, hydrogen electrode, buret, and stirrer. A small test tube is cut off a t the bottom and a 2-hole stopper carries the hydrogen electrode and hydrogen gas intake. The tube is adjusted so that the open end dips beneath the surface of the solution in the flask. The electrode is then adjusted so that it projects beneath the tube just far enough to still make a contact with the solution when a bubble of gas is escaping from the tube, thus giving all the advantages of a rocking electrode with none of its disadvantages. Received August 25, 1922. Contribution No 7 from the Research Laboratory of the National Biscuit Company. a J . Am. Chem. Sor , 35 (1913), 847.
FIG.1
February, 1923
INDUSTRIAL AND ,VNGINE&RING CIfEMISTRXi
188
bridge have been demonstrated by Fales and Mudge3 io be the KYDPOCHLORIC, SULFURIC, AKD CITRICACIDSWITH SODIUM most reliable. They give constant results over long periods of HYDROXIDE time. Fales and Vosburgh' have shown that the contact potential differences between saturated potassium chloride and molar hydrochloric acid is zero, and in many other combinations it The titration curves for these three acids are given in is a t a minimum. Accordingly,the saturated potassium chloride- Figs. 3, 4, and 5, respectively. Both the dibasic sulfuric calomel electrode and acid and the tribasic saturated potassium chloride salt bridge citric acid give curves were used, wooden similar to the monopegs being employed basic hydrochloric to reduce diffusion. This gave results con- acid in that there is stant to 0.2 millivolt, only one inflection, care being taken not but citric acid is t o have the plugs tight much less active than enough to affect the either of the other voltage of the cell.' The mercury used was two and its endpurified by distilling point is not nearly under reduced pres- as sharp. While a sure.6 The calomel choice of indicators was prepared by the changing over a wide electrolytic method of Ellis.' The potassium range of pH is poschloride was purified sible in titrating hyby twice recrystallizing from distilled water drochloric and sul' I Baker's analyzed po- furic acids, great care tassium chloride and must be exercised in 0310 0 5 1.0 15 2.0 25 0302 M35NaOH 4.0 41 5.0 FIG. 2-oSSERV5D E. M. F . N PH A S then fusing in plati- selecting an indicator . DETERMINEDBY T€iE SYSTEM: num. Hgl HgCl KC1 (SATD.)/KC~(SATU.)IUNKNOWN HYDROGEN -C o mEI SOLN. I HrPt pressed hydrogen was Ez ELECTRODB used, purified by pas- when titrating citric N ~ O HBY MEANSOR HYDROGEN Et E El OBSERVED E. M. F. sing it successively acid. AT 26O C. Et = 0.5266 VOLT,ASSUMING Hg I HgCI.01 M through alkaline potasKC1 at 26'C. 0.5618 VOLT sium permanganate, The electrometric titrations of hydrochloric and sulfuric alkaline pyrogallate, cotton wool, and finally through distilled acids check the indicator titrations very closely, as is seen from water to saturate the gas with water and prevent removal of the figures. The citric acid was purified by repeated rewater from the titration cell. crystallizations and dried a t 103' C. Calculations on the CONDUCTIVITY WATER-The so-called "conductivity" water used in this work was prepared by distilling first from a sul- basis of 1 molecule of water of crystallization show an error of furic acid-dichromate solution and then from barium hydroxide. 2.4 per cent in the electrometric titration. Considering the Only a part of the shape of the infleetion of the curve for citric acid, this error is steam distilling over not excessive and is no greater than the error encountered was condensed. using indicators.10 E. M. F. MEASURING I N S T R U M E N T Owing to the nature of PRIMARYCALCIUM 09' some of the work in PHOSPHATEWITH o 8 this laboratory, certain cells of high inSODIUM HYDROXIDE ternal resistance, solutions of low concentraFigs. 6 and 7 show O 7 tion of electrolytes, have to be measured. the results of titrat- @. Accordingly, the ap- ing commercial pri- O paratus described by mary;calcium phosBeans and Oakes7 was installed and was used phate, making up 0 for this work. Any the solutions in boiled 0 high-grade instrument distilled water and 0 accurate to 0.1 millivolt is satisfactory for again in conductivity 0 4 Taking the 043 titration work. The water. method of calculating middle point of the 039 results is shown in inflection in each 0 3 5 FIG.3-TITRATION OF 30 CC. 0.1 M HCI WITH 0.1 M NaOH BY M ~ A NOF S HYDROGEN Fig. 2, which includes a graph for rapid con- curve, it is seen that ELECTRODB AT 2 5 O C. END-POINTFROM versions of observed the former requires CURVE 30.0 Cc. 0.1 M H C l a 29.86 Cc. 0.1 M 22-75 cc, of alkali corFIG. 5-TITRATION OB 0.2498 G. CITRIC NaOH S R ~ M TITRATION(USIIG PHENOL- e. m. f. readings in volts directly into pH responding to 89-25 ACIDIN 200 CC. CONDUCTIVITY WATER WITH PHTHALEIN) 30.0 Cc. 0.1 M HC1 * 29.99 Cc. 0.1 M NaOH BY MEANSOF HYDROGEN ELECvalues. (See Lewis, 0.1 M NaOH AT 250 c. Brighton and Sebas- per cent primary cal~ and cium phosphate in the tian,* Lamb and Larson,g Fales and V ~ s b u r g h ,Fales Mudge.*} sample, while the latter requires 22.00 cc., showing the sample
350/&44!4d
-
3
'
A m . Chem. SOC.,48 (1920), 2434. 4 Ibzd., 40 (1917), 191. Hulett, 2.physik, Chem,, 83 (1900), 641. J . pm. Chem. SOC., 38 (1916), 737. I b i d , , 49 (1920), 2116. IZbid., 89 (1917), 2245. s Zbld., 41 (IgZO), 229. 8.7.
to contain 86.80 per cent primary calcium phosphate. Gravimetric analysis of this sample showed it to contain 87.00 per cent primary calcium phosphate. Failure to completely of Columbia University, for s carried out in the authors'
INDUSTRIAL AND ENGINEERING CHEMISTRY
184
Vol. 15, No. 2
cc.
FIG.7-TITRATION OR 0.2965 G.Ca(HiP0dr (COMMERCIAL) IN 200 Cc. CONDUCTIVITY WATERWITH 0. 1 M NaOH BY MEANSOF HYDROQEN ELECTRODE AT 26' c.
FIG.6-TITRATION OB 0.2983 G.Ca(HnPO4h I N 200 CC. BOILED DISTILLED (COMMERCIAL) OF WATER WITH 0.1 M NaOH BY MEANS HYDROGENELECTRODE A T 25' c.
0 2 M NaOH
FIG. &-TITRATION OF 40 CC. 0.0711 M HsPO4, MADEUP TO 200 CC. WITH CONDUCTIVITY WATER,AND TITRATEDWITH 0.2 M NaOH BY MEANSOP HYDROGEN ELECTRODE A T 26' c.
0950 10910
0 830 0790 0 750 0710 0 670 0.630
0 590 0 550
0 510 0430 0
3
0350 1
[ 0
9
o 10
~
~
cc
OlMH P
15
20
'
2 5 '3P043!
!
I
~
/
4!0 4'5 5!0
F~Q. 9-TITRATION OB 40 CC. 0.1 M NaOH, MADE UP TO 200 Cc. WITH CONDUCTIVITY WATER,AND TITRATED WITH 0.1 M HsPOi BY MEANSOF HYDROGEN ELECTRODE AT 25'C.
~
I
'
~
FIG. IO-TITRATION OF 40 CC. 0.0888 M HIPOI, MADEU P TO 200 CC. WITH CONDUCWATER,AND TITRATED WITH 0.1 M Nap;COI BY MEANSOF HYDROGEN ELECTRODE AT 25' c. TIVITY
remove carbon dioxide from the water before making up the solution resulted in a false titration curve (Fig. 6), with an error of over 2 per cent in the final results. The great difficulty experienced by analysts in titrating solutions of primary calcium phosphate is readily explained by the curve of Fig. 7 . The indicators commonly employed for this work do not change at the correct pH and the inflection is not sharp enough for end-points. This coupled with the precipitation of secondary calcium phosphate makes the ordinary indicator titration of little value. ORTHOPHOSPHORIC ACID WITH SODIUM HYDROXIDE Orthophosphoric acid was prepared by refluxing phosphorus pentoxide several hours and allowing it to stand two days before standardizing: The results of titrating this acid with sodium hydroxide are shown in Fig. 8. The chief points of interest are that there is no third inflection due to the third hydrogen ion of the acid, and the second end-point requires more than twice as much acid as does the first. Several other titration curves were made on various
FIG.I~-??ITRATION OF 0.1734 G. NarCO,: IN200 CC.
BOILEDDISTILLEDH i 0 WITH 0.1 dzf HsPO4 BY MEANSon HYDROGEN ELECTRODE n~25'C.
samples of phosphoric acid, and both these points were confirmed in every case. In a titration of a portion of this same acid using color indicators the methyl-orange end-point required 14.20 cc. of alkali, thus checking very closely 14.25cc. as determined electrometrically. The phenolphthalein end-point required 28.53 cc. alkali as against 29.35 cc. from the curve. It is evident, then, that through the influence of the third hydrogen ion of the acid the second inflection is inhibited a t its start and prolonged at the end, thus raising the whole section of the curve. Fortunately, phenolphthalein changes color at a point still on the steep part of the inflection and where exactly twice as much alkali is used up as for the first endpoint. l1 Reversing the operation and titrating sodium hydroxide with phosphoric acid does not give an exact mirror image of Fig. 8. Such a reverse curve is shown in Fig. 9. The 1 ' The presence of primary phosphate would cause the second endpoint t o require more than twice as much alkali as the first, but every preIt is certain that Fig. 8 reprecaution was taken t o eliminate this error. sents the facts for o-phosphoric acid.
of the same acid required 16.21cc. of the same carbonate soh-
IN
?h 13-TITRATION . OF 0.5365 G. NazCOi 200 Cc. BOILED DISTILLEDWATERWITH
0.2 M HCI BY MEANSOF HYDROGEN ELECTRODB h T
25' c .
First
ZNaaCOa 3. H3POI = 2NaHCOa Second end-point :ZNaHCOa Po4 2Ha0 -I- 2CO2.
+
ence in the character of the curve. Three times as much phosphoric acid (30.80 cc.) are required to reach the first endpoint (7.7 cc.). A study of the conditions necessary for equilibrium a t these two points shows that the following equations represent the facts, and also shows why it requires three times as much acid to titrate from the first to the second end-point as to titrate to the first.
+ Na2HP04.
+ NazHPOa f 3H&'04
end-point: = 4NaH2-
These equations were checked by an indicator titration of
the same as indicated
'
'
!c
02MHCI
I
1 7
with sodium carbonate, the same type of curve is obtained as for the second portion of Fig. 10. The inflection is less abrupt as shown in Fig. 12, and the end-point can only be guessed. SODIUMCARBONATE AND HYDROCHLORIC ACID The sodium carbonate titration is particularly interesting because of its application to standardizing solutions. By titrating the solution rapidly the curve represented in Fig. 13 is obtained. The first inflection is not very abrupt and judgment of the end-point is therefore correspondingly inaccurate. The second inflection point is abrupt and steep, with the end-point well defined. The titration required about 4 hrs. for its completion. Fig. 14 represents the results of a duplicate titration, except that after the first endpoint was passed the solution was allowed to stand for some time with the hydrogen bubbling through it to remove the liberated carbon dioxide. This brought the curve back up to the true end-point a t pH 8.10, and of course gave a different shape to the second inflection. The second end-point was chosen at a point requiring just twice as much acid as the first. The DH 4.12 of this Doint. it will be noted, is the same as for the-lorresponding point in Fig. 13. These two curves show the nature of the end-points under different circumstances. The advantage of the methyl-orange over the phenolphthalein titration is obvious. 12
Clark, "Determination of Hydrogen Ions.'!
