July 15, 1929
INDUSTRIAL AND ENGINEERING CHEMISTRY
125
closed and the water level read as accurately as possible, using a reading glass if necessary. Example Apparatus tested-a laboratory wash bottle Barometric pressure = 729.5 mm. a t 23’ C.
31.039 cc. X 0.73554 (factor for cc. HzO to mm. Hg) = 22.830 mm. He 72657 - 229830 = 703.74 mm. pressure in flask (22.830 + 726.57)lOO = 3.1422 per cent change in pressure (36.760 cc. X 676.683) + 703.74 mm. = 35.347 cc. dry air in ~~~
=
726.57 mm. a t
0” c. Room temperature = 27.35” C. = 27.057 mm. saturated vapor
tension Water temperature = 27.35’ C. Vapor pressure in room air = 9.73 mm. Volume of space above 0.0 cc. mark = 8.978 cc. Buret reading = 36.69 cc. Height of water column = 31.513 cm. Correction for capillarity = 0.33 cm. water The corrected volume read on the buret is the total expansion of gases from all sources, and the corrected volume of the column of water is the pressure change corresponding to the change in volume. This increase in volume includes the expansion of the gas mixture in the flask, expanding as a true gas since no water was present, the expansion of the 8.978 cc. of air above the water in the buret, and the vapor added in the buret to change the mixture from a state of partly to one of completely saturated vapor. Calculations I n computing the results the best procedure is first to eliminate both the amount of vapor added in the buret and the total originally saturated volume above the buret water, and in the example given that plan is followed: 36.69 cc. x 1.0019 (buret factor) = 36.760 grams or cubic centimeters corrected
total expansion (8.978 cc. X 699.513) f 726.57 mm. = 8.644 cc. dry air in original saturated volume (8.644 cc. X 726.57) c 703.74 mm. = 8.925 cc. dry air a t 703.74 mm. Hg pressure 8.925 - 8.644 = 0.281 cc. expansion 35.347 cc. (dry air in total expansion) - 0.281 cc. = 35.066 cc. dry air from flask in volume increase At this point the vapor which was added and the 8.978 cc. originally over water have been completely eliminated from the problem. The original vapor pressure in the flask was 9.73 mm., which is equal to 1.339 per cent of the total pressure and hence the dry air = 98.661 per cent of the total. (35.066 c 98.661)lOO = 35.542 cc., which represents 3.1422 per cent of the total volume. (35.542 + 3.1422)lOO = 1131.12 cc. total, and subtracting from this the 35.54 cc. increase leaves 1095.68 cc. as the volume of the flask. The theoretical volume computed from the weight of water required to fill the flask was 1096.25 cc. Trials with two flasks coupled together gave on one day a value of 2143.78 cc., and on another day 2144.53 cc. The theoretical volume for the two flasks was 2144.65 cc. These results are considered satisfactory, and show the accuracy of the method when operated under fairly good conditions. Literature Cited (1)
Chemiker Kalender, 1923, Pt. 11, p. 127.
Bimetallic Electrodes for Titrations Involving a Change of Hydrogen-Ion Concentration’ Raymond Matthew Fuoss SKINNER, SHERMAN & ESSELEN, INC., BOSTON, MASS.
The following electrode pairs have been used for POmotor stirrer, buret, and a HE necessity arose in tentiometric titrations involving a change of hydrogenpair of electrodes, which were this laboratory for a method of t i t r a t i n g ion concentration: antimony-lead, antimony-amalconnected to the potenticertain highly colored colgamated copper, bismuth-silver, and copper-copper ometer through a reversing loidal a l k a l i n e s o l u t i o n s . oxide. All of these gave satisfactory results if due switch. The electrodes were Bimetallic electrode systems precautions were taken; for general acidimetric a n d soldered to copper lead-wires, for oxidation-reduction titraalkalimetric work, however, the writer prefers the and were simply hung Over antimony-amalgamated copper Pair. the edge of the beaker so that tions have been described by The use of two suitably chosen metallic electrodes they were about half im. Willard and Fenwick (8)* and by firman ( 3 , 4 , 5 ) . Furman as discussed in this paper suggests various industrial mersed. The soldered junc(4) a l s o s t a t e s t h a t t h e applications-for example, indicating, recording, and tion was coated with paraffin amalgamated gold electrode Process control equipment in cases where a change of to prevent contamination of may beused to followneutraliPH occurs.z The applicability of the method for the electrode when the eleczations of hydrochloric acid volumetric determinations of metals has been sugtrode was rinsed with acid. or sulfuric acid, and Todd gested. Method (7) has patented the use of A definite volume of “unknown” was pipetted into the the cadmium-antimony alloy-iron pair for determination of hydrogen-ion concentrations. It was therefore decided to beaker and diluted to about 150 cc. The electrodes were investigate further the possibilities of such systems for the inserted and’ standard solution was added a t one-minute problem in question. The characteristics of four electrode intervals in successive portions of decreasing size until the pairs have been studied and are reported in this communica- magnitude of the change of voltage with portion of solution added gave warning that the end point was near. The tition. tration was then carried over the end point by adding at 60Apparatus second intervals equal portions (usually 3 drops) of standard The apparatus consisted of a Lee& and Northrup No. 7654 solution. The electromotive force usually became nearly pH indicator (one scale division equals 2 mv.), titration beaker, but not quite constant in 60 seconds; the value a t exact 60-second intervals was noted for the curve, however. The 1 Received November 3, 1928. reason for this procedure is twofold: First, the location of * Patent application has been made to cover such industrial uses. * Italic numbers in parenthesis refer to literature cited at end of article. the end point on the graph is made sharper, as a rule, by this
T
ANALYTICAL EDITION
126
AE/AV
600
400 0
CP EP EP Figure I-Titration of S o d i u m Hydroxide w i t h 0.1 N Hydrochloric Acid Abscissas marked EP are calculated end points. Abscissa units are cubic centimeters.
method for antimony-amalgamated copper and for antimonylead, because the drift is in opposite directions on opposite sides of the end point, as will be pointed out later; and second, the total time of titration is shortened by adding portions a t one-minute intervals rather than waiting until the e. m. f. is steady after each portion, which would require perhaps 10 minutes per portion, as indicated by experiments made in buffer solutions which will be discussed later. The drift was greatest in titrations of sodium hydroxide or hydrochloric acid, weaker with acetic acid, and almost negligible with sodium -8 carbonate. The value of AE/ AV was then plotted against the cubic centimeters of titer solution, and the end point read from the graph. The end point can also be estimated quite closely by inspection from t h e values of the e. m. f. as noted. Antimony-Lead
The antimony electrode (5 X 1 X 0.5 cm.) was made from a stick of cast antimony by pour+ ing molten chemically Boo p u r e antimony into a glass t u b e held in a 400 p l a s t e r of Paris cast. The bar was filed until O all surface pits were reeo LC m o v e d a n d was then Figure 2-Titration of Hydrochloric rubbed with emery paper Acid w i t h 0.1 N Sodium Hydroxide Abscissas marked EP are calculated end until smooth* It "as I points. c l e a n e d w i t h concenAbscissa units are 0.5 cc. trated hydrochloric acid and rinsed with distilled water before each titration. The lead electrode was a cylinder 6 em. long and 0.8 cm. in diameter, cut from a bar of commercial lead. It was cleaned before each titration with 6 N nitric acid, and rinsed with distilIed water. I n several of the titrations indicators were added and their I P
IC
VOl. 1, No. 3
end points noted. The indicators used were phenolphthalein (8.3 to 10.0), bromothymol blue (6.0 to 7.6), and methyl orange (3.1 to 4.4). These colorimetric end points (1) are indicated on the titration curves A (P = phenolphthalein, MO = methyl orange, BTB = bromothymol blue, etc.) in Figures 1 to 3. It will be noted that a characteristic break is obtained a t the bromothymol blue end point, which corresponds to exact neutrality. The electrometric end points also check within the experimental error the stoichiometric end points, which are indicated by the abscissas marked EP on Figures 1and 2, as calculated from the volumes of solutions used. (The normalities of the standard solutions were determined by the usual methods.) Curve A in Figure l is for sodium hydroxide titrated with 0.1 N hydrochloric acid; in Figure 2, for hydrochloric acid with 0.1 N sodium hydroxide; and in Figure 3, curve A shows the two end points obtained by titratirlg sodium carbonate with 0.1 N hydrochloric acid. The electrometric end points are tabulated on the figure. The plots of AE/AV are also given on the corresponding figures. In these experiments the maximum possible accuracy was not attained, since the entire titration was carried out with tenth-normal solutions, where the accuracy was a t best 0.05 cc., or 0.2 per cent for a 25-cc. titration. f 260
400
240
iao
BO
300
no
of S o d i u m Carbonate w i t h 0.1 N Hydrochloric Acid Ordinates for A , upper left; for B, lower left; for C, upper right.
Figure 3-Titration
The value of reading the e. m. f. before it reaches constancy is illustrated by Figure 4. Before the end point is reached the e. m. f. tends to fall with time; that is, the noted values (dotted curve) are higher than the values ultimately reached (solid curve). Beyond the end point the e. m. f. tends to rise, and hence the noted points are lower than the constant vaIues. Therefore, three inflections in the curve will appear, and the middle one, which corresponds to the end point, will be marked by a minimum between two maxima in the aE/AV curve. The approach of the end point is also clearly marked. During the addition of the first 90 to 95 per cent of the acid the curve slopes down; it then bends upward more and more sharply as the end point is approached. The chief precaution to be taken with this pair of electrodes is to be certain that no oxide or carbonate adheres to the lead electrode, which may consume extra acid if the solution comes in contact with it. The electrodes are apparently slightly attacked by the solutions, and the end point of the first titration is usually t o b low. With this approximate end point as a guide, however, a second more accurate titration can be made by adding about 95 per cent of the necessary acid thus determined before inserting the electrodes.
INDUSTRIAL AND ENGINEERING CHEMISTRY
July 15, 1929
127
Table I-Titration of Acetic Acid with S o d i u m Hydroxide (25.0 cc. of 0.1124 N acetic acid diluted to 150 cc. and titrated with 0.1 N NaOH) cc. 0.00 5.0 10.15 15.02 20.15 25.20 27.10 27.50 27.80 27.95 28.08 28.20 28.30 28.50 -E 154 182 190 197 203 223 240 246 256 264 269 286 302 309 AE/AV 5.6 1.6 1.4 1.2 4.0 8.9 15 33 53 38 142 160 35 9 End point = 28.25 cc. Calculated end point = 28.10 cc.
Bismuth-Silver
The bismuth electrode (3 X 0.5 X 0.5 cm.) was filed from a rough casting made of commercial bismuth. The silver electrode was a 1.5-mm. pure silver wire, immersed to a depth of about 2 cm. The bismuth electrode was cleaned with concentrated h y d r o c h l ori c acid before use, the silver electrode with dilute nitric acid followed by ammonia. Curve B in Figure 1 is for the titration of caustic with hydrochloric acid, B and B’ in Figure 2 are for hydrochloric acid with sodium hydroxide, and B in Figure 3 is for carbonate with hydrochloric a c i d . The electromotive force of this pair became practically constant within 60 seconds when acid or carbonate was titrated, and Figure 4-Comparison of Titration the end point of bromoCurve Obtained a t Equilibrium (Solid) w i t h T h a t Obtained at One- thymol blue was found to M i n u t e Intervals (Dotted) correspond to the inflection point of the curve for the titration of hydrochloric acid. It will be noted (B and B’ in Figure 2) that sometimes the break near the end point was about 100 mv. and a t other times 200 mv. It was found that the size of this break depended on the pretreatment of the silver wire. Heating i t in the gas flame and allowing to cool, treating with nitric acid, or using as the anode for a minute in a potassium sulfate solution gave an electrode which would give a 200-mv. break. Washing with ammonia or stannous chloride or using as the cathode for a short electrolysis gave an electrode which would give a 100-mv. break. It appears as if the larger -€
I
I
I
I
I
1
290
-E 210
320
I30
24 0
I60 0
20
30rc.
Figure 5-Titration of Acetic Acid with 0.1 N S o d i u m Hydroxide Using Antimony-Amalgamated Copper Pair Ordinates for A and B , upper left; for C, lower right.
break is caused by the silver action as a silver oxide electrode, and the smaller break by the wire action as a silver electrode, or more probably a silver chloride electrode since chloride ion is present during the titration. As was the case with the antimony-lead pair, the accuracy of the titration is increased by inserting the electrodes after
28.95 313
most of the titer has been added. This pair of electrodes gave the most concordant results when used for titration of hydrochloric acid; when used for the titration of sodium hydroxide, however, the results were quite erratic, and the end point frequently difficult to locate. This pair is therefore not recommended for titration of strong bases. Antimony-Amalgamated Copper
The amalgamated-copper electrode was prepared by dipping a copper wire into a saturated mercurous nitrate solution. It was redipped before each titration. The amalgamated-copper electrode was positive with respect to the antimony electrode. Curve C in Figure 1 shows the titration of sodium hydroxide with 0.1 N hydrochloric acid, in Figure 2, of hydrochloric acid with 0.1 N sodium hydroxide, in Figure 3, sodium carbonate with 0.1 N hydrochloric acid, and in Figure 5, acetic acid with 0.1 N sodium hydroxide. The data for the last curve are given in Table I. It will be noted that curve C in Figures 1and 2 goes through three inflection points, of which the middle one is located a t the bromothymol blue end point, as was the case with the antimony-lead pair. The tendency is so pronounced here that the curve actually goes through a maximum and a minimum before continuing, which locates the end point very sharply. Curves A and B of Figure 5 show that the same thing occurs with the titration of acetic acid, although the break is smaller; curve C was obtained by allowing the e. m. f. to become practically steady before a further portion of alkali was added. The difference between curves A and C of Figure 5 is the same as that between the dotted and solid curves of Figure 4. Copper-Copper Oxide
The copper oxide electrode was prepared by using a copper wire as the anode for several minutes in a dilute potassium ferricyanide solution, the cell being connected in series with a 10-watt lamp to the 110-volt circuit. A smooth, adherent, chocolate-brown coating forms in a few minutes. The nature of this deposit was not investigated in detail owing to lack of time. This particular electrode was chosen after a series of experiments had been made in an effort to find one suitable for the titration of copper, since many metals could not safely be inserted in a copper solution and this electrode was found to be satisfactory for the purpose. The coating contains copper and iron; and a copper wire, blackened in the flame, gives the same e. m. f . in an alkaline solution against,a bright copper wire as the electrolytically prepared electrode, but the flame-prepared electrode rapidly shifts its potential, while that of the electrolytically prepared electrode remains constant, Although the coating contains iron, therefore, this electrode will be called for convenience a copper oxide electrode. A bright copper wire was used as the other electrode. Figure 6 and Table I1 give the results of the titrations of 0.1 copper sulfate solution with 0.1 N sodium hydroxide solution. Table 11-Titration of Copper Sulfate with S o d i u m Hydroxide 0 986 M CuSOi 0.1 N NaOH NaOH Found Calcd. CC.
CC.
CC.
5.0 10.0 25.0
7.50 14.75 37.00
7.4 14.8 37.0
ANALYTICAL EDITION
128
Vel. 1, No. 3
The end point, as characterized by the sharp break in the curve, appears when 3 mols of sodium hydroxide have been added to 2 mol9 of copper sulfate. This corresponds to the precipitation of all the copper according to the equation: 4CuSO4
+ 6NaOH .--f 3Cu(OH)p.CuSO4 + 3Na2SO4
The solution a t the end point is neutral to bromothymol blue, and contains no copper detectable by ferrocyanide or ammonia. -E
Figure 7-Change of e . m . f . between Antimony-Amalgamated Copper w i t h p H OP Solution
/so
/a0
50
0 cc.
io
20
30
40
solutions of varying hydrogen-ion concentration. The slope, dE/d(pH) is 0.042, which is lower than the value 0.059 obtained by Roberts and Fenwick (6) for the cell Sb I Sb&, solution 11 3.5 N KC1, AgCl 1 Ag a t 25’ C., and the value 0.054 obtained by Franke and Willaman (2) for the cell Sb I solution I/ 1.0 N KC1 1 Hg2C12 I Hg. This lower value of the slope may be due to incomplete equilibrium between electrodes and solution, or more probably, to variation of the amalgamated copper electrode with pH in the same direction as the antimony electrode, although to a much smaller extent. This was checked by determining the e. m. f. of the amalgamated electrode against the calomel electrode when the former was placed in different buffer solutions. Acknowledgment
Figure 6-Titration of Copper Sulfate Solution w i t h 0.1 N Sodium Hydroxide Solution
Discussion This group of experiments indicates the possibilities of this method in determining the composition of precipitated hydroxides and basic salts, and also indicates that this general method can be used to follow any titration involving a change of hydrogen-ion concentration. From the data presented it is apparent that titrations involving a change of hydrogen-ion concentration may be followed potentiometrically by measuring the e. m. f. between two suitably chosen metal electrodes inserted in the titration mixture. Figure 7 gives the values of the e. m. f. between the antimony-amalgamated copper pair when placed in buffer
The writer takes this opportunity to thank G . J.Esselen, Jr., of this organization, and G. S. Forbes, of Harvard University, for their criticisms of this manuscript and of the results presented therein. Literature Cited (1) Clark, “Determination of Hydrogen Ions,” pp. 78 and 80, Williams and Wilkins Co., Baltimore, 1925. (2) Franke and Willaman, IND.END.CHBM.,20, 87 (1928). (3) Furman, J . Am. Chem. Soc., SO, 268 (1928). (4) Furman, I b i d . , SO, 273 (1928). (5) Furman and Wilson, I b i d . , SO, 277 (1928). (6) Roberts and Fenwick, I b i d . , 50, 2137 (1928). (7) Todd, U. S . Patent 1,601,383 (September 28, 1926). (8) Willard and Fenwick, J Am. Chem. Soc., 44, 2504 (1922).
Quantitative Determination of Formaldehyde in a Pharmaceutical Preparation’ Oscar Heim 244
w
EAST81ST STREET, NEW Y O R E ,
HEM called upon to determine formaldehyde quantitatively in a pharmaceutical preparation, the writer found the well-known methods-hydrogen peroxide, iodine, and ammonia methods-to give unsatisfactory results, due probably to the presence of a large number of different substances, both organic and inorganic. He therefore devised the method described below. The results by this method as compared with those by other methods show it to be very satisfactory: iodine 0.019 per cent, peroxide 0.63 per cent, silver chloride (present method) 0.20 per cent. The actual content was 0.20 per cent. This method is not applicable in the presence of sugars. METHOD-TOa 10-cc. sample (if about 0.2 per cent formal1 Received
March 11, 1929.
N. Y.
dehyde is present) add 2 cc. of concentrated hydrochloric acid and 10 cc. 1 M silver nitrate. Shake and add immediately 4 cc. of 50 per cent sodium hydroxide. Shake again and let stand for 15 to 30 minutes with occasional shaking. The presence of formaldehyde turns the mixture immediately black. Filter and wash with hot water. Perforate the filter by means of a thin stirring rod and rinse the filter with 1:3 nitric acid to dissolve all the reduced silver, leaving the excess of the silver chloride undissolved,’ After some dilution filter off the silver chloride and stir the filtrate with sufficient hydrochloric acid to precipitate the silver, which is determined in the usual manner. 2AgCl = 1CHgO
