Analysis of Boron Tribromide and Its Addition Compounds WILLIAM J. SCHUELEI, J. FRED HAZEL, and WALLACE M. MCNABB Department of Chemistry, University o f Pennsylvania, Philadelphia, Pa.
Several methods of determining the boron and bromine content of boron tribromide and its addition compounds w-ere compared. The sample was hydrolyzed and the organic portion removed by means of an ion exchange resin to yield an effluent containing boric and hydrobromic acids. The inorganic acids were titrated with sodium hydroxide by four different methods: potentiometric, conductometric, visual, and spectrophotometric. Various acid-base indicators were employed. The potentiometric and spectrophotometric methods proved to be equally accurate, but the spectrophotonietric titration was faster.
Table I.
Comparative Titrations by Different Methods
( B r / B Ratios Given) SpectrophotoPotentioVisual metric metric JIethyl Red and Phenolphthalein 3.604 3.604 3.482 3.604 3,578 3.632 3.594 3.604 3.604 3.GO4 3.604 3.456 3.606 3.6QQ
3.688 3.545 4.000 3.6OG
hlixed Indicator and Phenolphthalein 3.488 3.601 3,474 3. 604 3.530 3,5i6 3.515 3.604 3.G10
A
N EXBAIINATION of the literature on the analysis of boron tribromide indicates that relatively poor accuracy is usually obtained in the boron analysis (3, 4, 6, 7'). Moreover, the analysis of boron tribromide addition compounds is further complicated by the fact that the amount of boron in the addition compound is usually only a small percentage of the total composition. The method most frequently used is hydrolysis of the salt, titration of the total acids (boric and hydrobromic) with standard sodium hydroxide, followed by a determination of the bromine content by titrating with silver nitrate. The boric acid is then obtained by difference. In the present work it was found that the basicity of the organic molecules used in preparing the addition compounds interfered with the determination of the acids produced upon hydrolysis. The titration values were inaccurate because of salt formation between the organic base and the hydrobromic acid. To overcome this difficulty, the hydrolysis mixture was passed through a cation exchange resin, Amberlite IR-100, hydrogen form. The organic base displaced its equivalent of hydrogen ion and was itself retained on the resin column. Some examples of the organic bases removed by this method are acridine, 5-nitroquinoline, morpholine, pyrrole, piperidine, 2- , 3- , and 4-picoline, isoquinoline, and nitrobenzene. The effluent containing the hydrobromic and boric acids v a s collected. The acids were quantitatively recovered in the effluent. A tenfold volume of nash water was found to be adequate to ensure that all the acids produced were transferred from the column to the collecting beaker. This amount of washing resulted in a solution approximately 0.02 to 0.0312' in acid. To evaluate various methods for determining the amount of hydrobromic and boric acid released, standard solutions m r e prepared and mixed to give an approximately 3.6 to 1 ratio of hydrobromic to boric acid-i.e., a ratio of bromine to boron of 3.6 to 1. These methods are compared in the body of the paper.
Conductometric
APPARATUS
A Leeds & Northrup p H meter, Type 7662, with glass and calomel electrodes and a magnetic stirrer for agitation, was used in potentiometric titrations. A Beckmann hIodel B spectrophotometer was adapted t o use a titration cell siniilar to that described by Smeetser and Bricker (8). A Rehberg microburet was adapted with a delivery tube t o permit titrations in the spectrophotometer cell. An Industrial Instruments conductivity bridge, Model R C lB, with platinized platinum electrodes and a magnetic stirrer n-as used in the conductometric titrations. COMPARATIVE TITRATIOXS
Potentiometric Titration. One hundred milliliters of the test solution mere titrated potentiometrically Tyith standard sodium hydroxide solution, using a glass electrode and a calomel cell. After the first end point (due to the neutralization of hydrobromic acid), approximately 5 grams of mannitol ( 6 ) mas added and the titration continued until all the boric acid was neutralized. For most of this n-ork and for general use, a 10-ml. microburet graduated in 0.02 ml. was used. However, to follow the instrument readings through the rapidly changing region of the end point, the titrations were taken to within 0.1 ml. of the end point, determined previously with a IO-ml. microburet. A Rehberg microburet was employed to add 0.005-ml. increments of the same strength base. The results are given in Table I. Visual Titration with Indicators. METHYLRED. Five drops of the methyl red indicator was added to 100 ml. of the test
t
REAGENTS
A 0.3N carbonate-free sodium hydroxide solution was prepared and standardized against standard hydrochloric acid. Standard solutions of about 0.015N hydrobromic acid and about 0.005N boric acid were prepared with carbon dioxide-free water. hIethyl red, methylene blue, and phenolphthalein indicators were prepared in the usual way. The mixed indicator ( 1 ) was prepared by mixing 40 ml. of methyl red with 15 ml. of methylene blue solution.
400
500 WAVE
600
700
750
Figure 1. Transmittance curves for mixed indicator A.
Methyl red-methylene
blue (acid)
B . Methyl red-methylene blue (basic) 1
.
LENGTH M p
C. Methyl red-methylene
Present address, The Franklin Institute, Philadelphia Pa.
(basic)
505
blue with plienolphthalein
and mannitol
ANALYTICAL CHEMISTRY
506
solution, which was then titrated with standard sodium hydroxi l e solution to the yellow end point. Approximately 5 grams of mannitol and 5 drops of phenolphthalein were added and the titration was continued to the pink color of phenolphthalein. Results obtained by this method are included in Table I. R ~ I X E D INDIC.4TOR. The titration Il-ith mixed indicator w : ~ carried out in the same manner as the methyl red titration, usir g 5 drops of mixed indicator instead of methyl red. The titration was carried to the gray-green end point, then, after the addition of mannitol and phenolphthalein, t o the red-violet color of the mixed indicator. The results are also given in Table I.
'+I
100
90
70 w
u
60
t
2
50
a iLII
s
Analysis of Boron Tribromide and Its Addition Compounds
(Spectrophotometric titrations x-ith mixed indicator and phenolphthalein) HBr, HJBOI, Br/B Coinpound Meq. Meq. Ratio Boron tribroiniden 1.551 0.517 3.000 -4ddition compounds b 5-Xitroquinoline 0,962 0.320 3.006 Quinoline 1.715 0.5757 2,979 4-Picoline 0.514 0.172 2.988 a hlandell ( 6 ) found a Br/B ratiq of 3.76 for boron tribromide. b Kinney and Mahoney ( 4 ) obtained a CI/B ratio of 4.38 for boron trichloride addition compounds.
ANALYSIS O F ADDITION CORIPOUNDS
Approximately 3 meq. of the boron tribromide addition compound were weighed out on an analytical balance and dissolved in 50 ml. of carbon dioxide-free distilled water, warming gently when necessary. The solution was then passed slowly through approximately 65 grams of the cationic exchange resin Amberlite IR-100 packed in a Jones reductor tube measuring 50 em. in length and 2 em. in diameter. The effluent and wash water were collected in a 500-ml. volumetric flask, and the column waa washed until a total volume of 500 ml. was obtained. Each of the titrations was done on 100-ml. nliquots of the effluent. The results of the mixed indicator spectrophotometric titration of boron tribromide and several of its addition compounds are given in Table 11. This procedure has also been used by Henry ( 2 ) in the analysis of boron tribromide addition compounds.
80
5I-
Table 11.
40
SUMMARY
30 20
5.80
.84
.88
.92
.96
6.00
ML. OF NaOH Figure 2. Spectrophotometric titration curve of hydrobromic acid in hydrobromic-boric acid mixture
Of the four methods investigated, the potentiometric and spectrophotometric methods gave consistently good results. Visual titrations gave consistently accurate values for hydrobromic acid but high values for boric acid. Although conductometric titrations did not give values as reproducible as the other instrumental methods, they gave approximately correct values and did not require the use of mannitol. To check these procedures further, several analysts used the various methods on aliquots of the test solution and the boron-
Methyl red-methylene blue mixed indicator
Spectrophotometric Titration. MIXED IXDICATOR. The absorption spectra of the mixed indicators in the test solution were determined (Figure 1, curve A ) , then the hydrobromic acid v a s titrated with a solution of sodium hydroxide and the absorption spectra were determined again (curve B ) . Mannitol (5 grams) and phenolphthalein ( 5 drops) were added, the boric acid was titrated, and a final absorption spectrum mas determined (curve C). From these three spectra the optimum \\*ave length for the titration was determined as 532.5 mp. At this wave length, in order to keep the readings on the scale at all times during the titration, the slit width was maintained a t 0.14 mm.; this may vary with the instrument and should be determined before the titration. \$'hen the apparatus was on these settings, the end point occurred a t 50% transmittance. The curve obtained in the spectrophotometric titration of hydrobromic acid with sodium hydroxide is shown in Figure 2. Figure 3 shows the titration of Boric acid with sodium hydroxide. A summary of the results is given in Table I. METHYLRED INDICATOR. By following the procedure used for the mixed indicator, the optimum wave length for the titration with methyl red was determined t o be 550 mg. The slit width mas 0.17 mm., and 50% transmittance corresponded to the end point. Spectrophotometric titrations of hydrobromic acid and of boric acid with sodium hydroxide mere made. The results are summarized in Table I. Conductance Titration. The platinized platinum electrodes of the conductivity bridge n-ere immersed in 100 ml. of the test solution and the resistance was measured. Increments of 0.5 ml. of standard sodium hydroxide solution mere added and the Fesistance was measured after each addition. Resistance readings were then converted to conductance values, corrected for volume changes, and plotted against milliliters of titrant added. The intersections of the straight lines were talien as end points. The results are given in Table I.
t a 2: 40 CT I-
$ 30 20
7.50
,52 ML.
,54
.56
.58
OF NaOH
Figure 3. Spectrophotometric titration curve of boric acid in hydrobromic-boric acid mixture Methyl red-methylene blue mixed indicator with phenolphthalein a n d mannitol
V O L U M E 28, N O . 4, A P R I L 1 9 5 6 bromide addition compounds. The largest disagreement was apparent in the visual titrations using only methyl red indicator; the mixed indicator visual titration mas only slightly better, Data obtained by the various analysts checked very well in both the spectrophotometric and potentiometric titrations. Conductometric titrations were not checked. While potentiometric and spectrophotometric methods are equally satisfactory, the spectrophotometric method appears to be more satisfactory for general use, because it is not necessary to plot the curve to obtain the end point. Once the apparatus has been standardized, the titration can be run to 50% transmittance and the buret reading taken. The needle deflection is also much greater on the spectrophotometer, making i t easier to take readings. Spectrophotometric titrations are especially suitable to visual titrations in which the end point is difficult to see.
507 LITERATURE CITED
Brunishols, G., Bonnet, J., H e h . C‘him. Acta 34, 2074 (1951). (2) Henry, 11. C., Ph.D. dissertation, University of Pennsylvania, (1)
1955.
(3) Johnson, H. It., J . Phys. Chem. 16, 3 (1912). (4) Iiinney, C. IT., ;\Ishoney, C. L., J . Org. Chem. 8, 526 (1943). (5) Kolthoff,I. AI., Sandell, E. B., “Textbook of Quantitative Inorganic Analysis,” 3rd ed., p. 534, Macmillan, Yew York, 1952. (6) RIandell, K. C., Ph.D. dissertation, University of Pennsylvania, 1951. (7) llartins, C. A , Ann. Chem. Justus Liebigs 109, 79-82 (1859). (5) Smeetser, P. B., Bricker, C. E., A x . 4 ~ CHEJI. . 25, 253 (1953). RECEIVED for review August 5 , 1955. Accepted January 18, 1956. Delaware Valley Regional Meeting, ACS, Philadelphia, Pa., February 16, 1956. Abstracted from a portion of a dissertation presented by William J. Bchuele to t h e Faculty of t h e Graduate School of the University of Pennsylvania in partial fulfillinent of the requirements for t h e degree of doctor of plulosophy.
Available Oxygen in Manganese Dioxide M. J. KATZ,
R.
C. CLARKE’, and W. F. N Y E
Signal Corps Engineering Laboratories, Fort Monmootb, N.
Examination of analyses of manganese dioxide for available oxygen indicated very poor agreement among different laboratories. An appraisal was made of the t w o most important methods used in this analysis-the ferrous sulfate method and the sodium oxalate method. It was discovered that good agreement between the methods could be achieved only with samples of natural origin. It was shown that in the analysis of synthetic m’anganese oxides, the oxalate method is subject to a positive error because of air oxidation of the oxalate. The methods gave good agreement when a protective atmosphere was provided in the oxalate procedure.
J. Exactly 6 grams of ferrous ammonium sulfate is introduced into the flask and a 0.5-gram sample of the manganese dioxide to be analyzed is added. Carbon dioxide is passed into the flask, the contents are brought to a boil, and thereaction is allowed to proceed until all of the sample is decomposed. Then 5 ml. of sirupy phosphoric acid is added and the excess ferrous ammonium sulfate is titrated with potassium permanganate. At the same time a blank is run with 6.000 grams of the ferrous ammonium sulfate under identical conditions. The available oxygen, calculated as per cent manganese dioxide, is given by the equation
where
TY = weight of sample B = volume of potassium permanganate required for blank
F
OR the past several years this laboratory has been concerned with the evaluation of manganese oxides for use as depolarizers in dry cells. As a part of this program, the analysis of the oxides for available oxygen is of importance. The ferrous sulfate method (6) was used in specifications defining the minimum available oxygen ( 4 ) . This procedure proved useful in establishing working hypotheses for the correlation of particle morphology and x-ray diffraction powder patterns with manganese to oxygen ratios. Difficulties arose, however, when interlaboratory comparisons were made. Even when the same methods of analysis were used, there were serious discrepancies in results. A survey of the literature showed this to be a problem of long standing. I n 1917, for example, Barnebey (1) pointed out that chemists could obtain discrepancies of 5% in the manganese dioxide content of a given sample, while smaller differences were not a t all uncommon. This has also been the authors’ recent experience. The ferrous sulfate and oxalate methods seemed t o offer most promise as potential referee methods. They are given by the National Bureau of Standards in their certificate for the analysis of manganese dioxide (standard sample No. 25b). Therefore, an investigation was undertaken regarding the general applicability of these methods to manganese oxides with different genetic histories. EXPERIMENTAL
Ferrous Sulfate Method. A 100-ml. portion of 10% (by volume) sulfuric acid is transferred to a 300-ml. Erlenmeyer flask.
* Present address, Chemistry N. Y.
Laboratory, 84th Police Precinct, Brooklyn,
T
N
= sample titer = normality of potassium permanganate
Sodium Oxalate Method. A 100-ml. portion of 10% (by volume) sulfuric acid and 1.000 gram of sodium oxalate are added to a 300-ml. Erlenmeyer flask, followed by a 0.5-gram sample of manganese dioxide for analysis. The reaction is allowed to proceed on the steam bath until the sample is decomposed. The excess oxalate is determined by titrating the hot solution with 0 . W potassium permanganate. Available oxygen as per cent manganese dioxide is calculated from the equation
% M n 0 2 = 4.346 (1/0.067 - T N ) TY
where the symbols have the same meaning as in Equation 1. DISCUSSION
Effect of Phase Type and Stoichiometry. For some samples, results obtained with both methods were consistent; for others there was considerable disagreement. I n the early stages of the work good agreement had been obtained for a number of naturally occurring pyrolusites. Of the various manganese oxides the pyrolusites @-manganese dioxide) make the closest approach to stoichiometric manganese dioxide. On the other hand, for other phase types usually having appreciable oxygen deficiencies, considerable discrepancies had been obtained. Phase type ( 2 ) means a commonly recurring manganese oxide, defined by its x-ray diffraction pattern and by a more or less distinctive morphology. Thus, it seemed that the phenomenon was related in some may to the manganese-oxygen ratio and/or the phase type. In order to test the dependence of analytical inconsistencies on stoichiometry and phase type, two materials, Samples D and 0, were prepared. Sample D is a synthetic P-manganese dioxide.
