ANALYTICAL
26
that the bismuth diethyldithiocarbamate complex is more stable than other metal diethyldithiocarbamate complexes. Anions. No interference was found from nitrate, sulfate, chloride, acetate, perchlorate, phosphate, tartrate, and citrate for the amounts of bismuth determined. Stability of Color. The yellowish coloration of the bismuth diethyldithiocarbamate complex in organic solvents—such as chloroform, carbon tetrachloride, and ethyl acetate—was found not to be very stable, especially with respect to light. The intensity of color gradually decreased with time of standing (Figure 3). When the carbon tetrachloride extract was allowed to stand at room temperature for more than 1 hour, it became turbid. Therefore, it is recommended that the absorbance be measured as quickly as possible and that the unknown sample be determined at the same time as the standard bismuth solutions. Downloaded via KAOHSIUNG MEDICAL UNIV on June 21, 2018 at 18:37:26 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.
DETERMINATION
OF BISMUTH
IN LEAD-BASE
was
one
selected.
Alloy.
Stability of Bismuth Complex
Bismuth, 0.20 mg. Bismuth, 0.04 O Wave length, 370 m^u Wave length, 400 m/u '
—
—
—
—
The lead-base bearing metal sample (NBS 53c,
con-
taining 10.20% antimony, 5.16% tin, 0.214% copper, 0.044%
arsenic, 0.0023% nickel, 0.0017% iron, and 0.093% bismuth) was treated in the following manner: A 1.0000-gram sample was dissolved in 20 ml. of 20% nitric acid by warming on the steam bath. After cooling, 3 grams of ethylenediaminetetraacetic acid and 10 grams of tartaric acid were added, followed by concentrated ammonium hydroxide (about 10 ml.) to adjust the solution to pH 7 to 8. The solution was then transferred to a 100-ml. volumetric flask and diluted to volume ivith water. The solution was sometimes cloudy. After being mixed thoroughly, an aliquot of 10 to 25 ml. of the solution was pipetted into a separatory funnel, followed by 10 ml. of water, 2 ml. of 5% sodium cyanide, 1 ml. of 0.2% sodium diethyldithiocarbamate, and 10 ml. of carbon tetrachloride. The mixture was shaken for 30 to 60 seconds, and the organic layer was filtered through a filter paper. The absorbance of the extract was measured at 400 µ, and the concentration of
bismuth manner
was read from a calibration curve obtained in a similar by the extraction of known amounts of bismuth.
The amount of bismuth in the alio}' was found to be 0.093 and 0.095%. These two values are in agreement with the average value of 0.093% as indicated by the National Bureau of Standards. LITERATURE
Chem., 23, 871 (1951). (4)
(5) (6) (7) (8)
Pribil, R., Collection
Czechoslov. Chem. Communs., 16, 542 (1953). Sandell, E. B., “Colorimetric Determination of Traces of Metals,” 2nd ed., p. 210, Interscience Publishers, New York, 1950. Scott, W. W., “Standard Methods of Chemical Analysis,” 5th ed., Yol. 1, pp. 150-61, X7an Nostrand, New York, 1939, Sedivec, V., and Vasák, V., Collection Czechoslov. Chem. Communs., 15, 260 (1950). Tompsett, S. L., Analyst, 63, 250 (1938).
Received for review May
HARVEY, JR., JOHN A. SMART, and EDWARD Department of Chemistry, University of Arkansas, Fayetteville, Ark. E.
When 1,10-phenanthroline is added to a solution containing both iron(II) and iron(III), a reddish orange iron(II) complex and a yellow iron(III) complex form immediately. The iron(II) complex has an absorbance maximum at 512 mg, at which wave length there is little absorption by the iron(III) complex. The two complexes have identical absorbance coefficients at A method is presented for the determination 396 mg. of iron(ll) and total iron in the same solution by simulof absorbance at 396 mg and at measurements taneous The limiting concentrations for interference 512 mg. by 22 cations and 14 anions are reported. 1,10-phenanthroline complex with iron(II) was first Walden, Hammett, and Chapman (4) used the complex as an internal indicator in the oxidimetric titration of iron with the ceric ion. A spectrophotometric determination of iron dependent on the formation of the iron(II)discovered by Blau (1). THE
CITED
(1) Cheng, K. L., and Bray, R. H., Anal. Chem., 25, 655 (1953). (2) Etheridge, A. T,, Analyst, 75, 279 (1950). H., and Wiberley, S. E., Anal. (3) LaCoste, R. J., Earing, .
12, 1954.
Simultaneous Spectrophotometric Determination of and Total Iron with 1,10-Phenanthroline AUBREY
per 10 ml. of CCh mg. per 10 ml. of CCI4
ALLOY
of the metals which commonly interfere with the determination of bismuth by other methods (2, 6). In order to test the reliability of the proposed method, the lead-base alloy Lead is
Figure. 3.
CHEMISTRY
S.
Accepted October
1, 1954.
Iron(ll)
AMIS
1,10-phenanthroline complex was developed by Fortune and Mellon (2). The iron(II)-l,10-phenanthroline complex, which is reddish orange in color, may be oxidized to a blue complex. Upon standing, this blue complex changes to a yellow complex, which also may be produced by complexing iron(III) and 1,10-phenanthroline directly. Harvey and Manning (3) showed that this yellow complex has a maximum absorbance at 360 mg and a small but measurable absorbance at 512 mg. Above 380 mg the absorption by 1,10-phenanthroline is negligible. In the method for the simultaneous determination of iron(II) and total iron reported in the present paper, advantage is taken of the difference in the absorption spectra of the reddish orange iron(II) and the yellow iron(III) complexes which are formed instantly on the addition of 1,10-phenanthroline to a solution containing these ions. EXPERIMENTAL
Instruments. A Beckman quartz spectrophotometer, Model DU, with 10-mm. silica or Corex absorption cells, was used for
VOLUME
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JANUARY
1955
spectrophotometric measurements. All measurements were made at maximum sensitivity of the instrument. A Beckman Model G pH meter was used to check the pH of all buffer solutions. Reagent Solutions. Reagent grade 1,10-phenanthroline monohydrate was obtained from G. Frederick Smith Chemical Co. The reagent was dissolved in distilled water which had been heated just to boiling. After it had cooled, the solution was diluted to the desired volume. A 0.1% reagent solution was used in forming the complexes for obtaining the complete absorption A 0.3% solution of the reagent assured a sufficient excess curves. in preparing the standard concentration curves and analyzing unknowns. Standard Iron Solutions. The standard iron(II) and iron(III) solutions were prepared by dissolving 7.0213 grams of reagent grade ferrous ammonium sulfate hexahydrate, (NH4)2Fe(S04)2.6H20, and 8.6337 grams of reagent grade ferric ammonium sulfate dodecahydrate, NH4Fe(S04)2.12H20, respectively, hi freshly distilled -water containing 3 nú. of concentrated sulfuric acid and diluting to a liter. The concentration of iron in each solution was determined gravimetrically as ferric oxide, Fe208, and was found to be 1.000 mg. per nú. The concentration of iron(II) in the former solution was determined periodically with ceric sulfate to avoid the use of the solution after any oxidation had occurred. Working standard solutions were prepared by dilution of these stock solutions. Buffer Solution. A 0.2M solution of potassium biphthalate gave a buffer of pH 3.98.
27
intersection of the absorption
curves of the iron(II) and iron(III) desirable. Two series of solutions were prepared containing, respectively, 6 p.p.m. of iron(II) and 6 p.p.m. of iron(III). The pH in each series was varied in intervals of 0.5 units from pH 2.8 to 5.5. Absorbance values for each solution were measured over the wave-length range from 390 to 399 mg.
complexes
was
Table I show's that the v'ave length of intersection remained constant at 396 mg for all solutions within the pH range from 3 to 5. The middle of this range, pH 4, was used in establishing standard concentration curves for the determination of iron(II) and total iron.
Table I.
Effect of pH
on
Wave Length of Intersection Wave Length of
Intersection,
µ
pH 2.8
395 396 396 396 396 396 397 398
3.0 3.5 4.0 4.5 5.0 5.2 5.5
Standard Concentration Curves. Solutions of the iron(II) complex and of the iron(III) complex were prepared by adding to aliquots of the desired stock solution, 10 ml. of 0.3% reagent solution, and 5 ml. of buffer solution, and diluting to 25 ml. with distilled w'ater. The concentration of iron in each of the two series of solutions ranged, in increments of 2 p.p.m., from 1 to 15 p.p.m. The absorbance of each solution was measured at both 396 and 512 mg. The absorbance at 512 mg of the iron(II) complex in concentrations greater than 10 p.p.m. vras too intense to be read in 10-mm. cells. The absorbance vs. concentration curves at 512 mg for iron(II) and iron(III), and a single concentration at 396 mg for both the iron(II) and iron(III) v'ere all curve linear over the concentration ranges measured. The slopes of these three lines vrere 0.196, 0.004, and 0.054 p.p.m.-1, respec-
tively. To establish the fact that the absorbances of the two complexes are additive at 396 mg, two other series of solutions v'ere prepared. In the first series, the total iron concentration was held constant at 10 p.p.m. The concentrations of iron(II) and of iron(III) were X and 10 X, respectively. The value of X was varied, in increments of 1 p.p.m., from 0 to 10 p.p.m. In the second series, the iron(II) concentration was held constant at 5 p.p.m. while the concentration of the iron(III) was varied from 0 to 10 p.p.m. in 1-p.p.m. increments. The absorbance of each of these solutions was read at 396 mg and at 512 mg. In the first series, the absorbance at 396 mg was constant and wras identical to the absorbance at 10 p.p.m. read from the standard concentration curve for 396 mg. The absorbance values at 512 mg, when corrected for the iron(III) interference at that wave length, duplicated the iron(II) concentration curve at 512 mg. In the second series, the absorbance values at 396 mg checked W'ith those from the corresponding standard concentration curve over the concentration range from 5 to 15 p.p.m., which was the range of total iron concentration. Moreover, the readings at 512 mg checked with the value from the standard concentration for iron(II) at 5 p.p.m. when suitable corrections were curve made for iron(III) interference. Standard curves at 512 mg should be prepared for both the iron(II) complex and the iron(III) complex. The standard solution of iron(II) used for this purpose may contain an excess of hydroxylamine hydrochloride. Because the absorbances are additive, the standard curve at 396 mg may be obtained at the same time from the solutions of either complex. Stability of Iron(III) Complex. After a short time the concentration of the iron(II) complex increased in solutions which contained both complexes with excess 1,10-phenanthroline. —
Figure 1. Absorption Curves of Iron—1,10-PhenanthroIine Complexes
Absorption Curves. The absorption curves for the iron(II) and for the iron(III) complexes were determined at several concentrations of iron between 2 and 10 p.p.m. Absorbance values were read at intervals of 5 mg in the wave-length range from 380 to 600 mg. All of the solutions for absorbance measurements were prepared by putting a suitable aliquot of the proper standard solution into a 25-ml. volumetric flask, adding 10 ml. of 0.1% reagent solution and 5 ml. of potassium biphthalate buffer, and diluting to the mark with distilled water. Figure 1 presents the absorption curves for the iron(II) and for the iron(III) complexes at iron concentrations of 3, 5, and 10 p.p.m. The iron(II) complex absorption curves show a maximum absorbance at 512 mg, at wrhich wave length there is little absorption by the iron(III) complex. The intersection at 396 mg of all curves for equal concentrations of iron(II) and iron(III) indicates identical absorbance coefficients for the two complexes
at this wave length. These curves suggest the possibility of determining iron(II) from the absorbance at 512 mg and total iron from the absorbance at 396 mg. Iron(III) then could be determined by difference.
Studies of the Effect of pH. The method of Fortune and Mellon (2) gives a range from pH 2 to 9 for the determination of iron(II). Harvey and Manning (S) showed that the iron(III) complex -was independent of pH over a range from 3 to 8. However, a study of the effect of change of pH on the wave length of
ANALYTICAL Table II.
Limiting Concentrations of Diverse Ions with 2
P.P.M. of
Iron(III)
Concn.,
P.P.M.
Cations Aluminum Ammonium
Arsenic(III) Cadmium
Calcium
Cobalt(II) Copper(II)
Lithium Magnesium
Manganese(II)
Mercury(I) Mercury(II)
Nickel Potassium Strontium Uranium
Zirconium(IV) Anions Acetate
500 500 100 50 25 500 250 10 5
10
1,000
500 500 15 10 1
2
1,000 500 100 50 100 2
500 1
Borate Bromide Carbonate
500 500 500 5
Chlorate Chloride Citrate
Fluoride Nitrate Nitrite
500
1,000
500 20 500 25
Oxalate
500 500 25 500
Sulfate
500
1
Apparent Concn. of Iron(III), P.P.M. 2.00 2.00 2.00 1.80 2.00 1.81
2.00 2.55
2.00 2.00 2.00 2.00 2.92
2.00 2.00 2.00 2.00 2.00 2.00
2.00 2.00 0.25 2.00 1.05 2.00 2.00 2.00 0.71 2.00 2.00 2.00 1.05 2.00 0.20 2.00
2.00 1.62 2.00
0.42
2.00 2.00
Even solutions of the iron(III) complex acquired a reddish tinge upon prolonged standing. To study the rate of the shift, two solutions were prepared. One contained 10 p.p.m. of complexed iron(III), and the other 5 p.p.m. of complexed iron(II) and 5 p.p.m. of complexed iron(III). The absorbances of both of these solutions were measured at 512 µ at frequent intervals over an extended period of time. No appreciable change in absorbance occurred for 30 minutes. After that time, the change in readings was marked. The increase in absorbance at 512 mg was greater for the mixture than for the iron(III) complex alone. After 15 days, the two solutions had approximately the same absorbance, indicating that an equilibrium between the two oxidation states had been reached. Further study of this shift will be undertaken. Interference of Diverse Ions with the Iron(III) Complex. A study was made of the effect of 22 cations and 14 anions on the The cations were absorbance of the iron(III) complex at 396 µ. present in solution as the chlorides or nitrates while the anion solutions were of the sodium or potassium salts. The concentration of iron(III) was held constant at 2 p.p.m. Aliquots of iron(III) solution and of the solution of the ion to be tested were placed in a 25-ml. volumetric flask, complexed with 10 ml. of 0.3% reagent solution, buffered with 5 ml. of buffer solution, and diluted to the mark with distilled water. The absorbance then was measured at 396 mg and the apparent concentration of iron The results are was read from the standard concentration curve. presented in Table II. The higher concentration of each diverse ion, where two concentrations are shown in Table II, or the single value given is the same as the limiting concentration found by Fortune and Mellon (2) in their study of interferences with the iron(II)-l, 10-phenanthroline complex. Barium and lead sulfates A precipitate were precipitated by the standard iron solution. formed in solutions containing molybdate ions when the complexing agent was added. Precipitation also occurred upon the addition of the buffer to solutions containing thorium(IV) ions. Reducing agents such as iodide, thiosulfate, and tin(II) reduced the iron to form the iron(II )-l, 10-phenanthroline complex.
CHEMISTRY
The limiting concentrations of cadmium, calcium, and cobalt with the iron(III) complex are one half of the values reported (2) with the iron(II) complex. Interference from manganese and from zirconium are considerably greater for the higher-valence complex. Because ferric ions complex more readily with anions
than do ferrous ions, the anion interference in many cases is more pronounced with the iron(III)-l,10-phenanthroline complex. Thus, the limiting concentrations for acetate, carbonate, citrate, fluoride, nitrite, and oxalate are much lower than the corresponding limiting concentrations with the lower-valence iron. For all other cations and anions tested, the limiting concentrations are at least as great as the concentrations permissible in the determination of iron(II), ANALYSIS
OF UNKNOWN SAMPLES
Solid mixtures of ferrous ammonium sulfate and ferric ammonium sulfate were prepared by weighing each salt and grinding them together in a mortar. The samples were sealed under an atmosphere of nitrogen to prevent oxidation of the iron(II), and were issued to one of the authors as unknowns. The proportions of iron(II) to iron(III) could not be calculated from the weights of the components because oxidation and loss of water of crystallization during grinding could not be prevented entirely. However, the concentrations of iron(II) and iron(III) in each sample were determined independently of the method under investigation as follow's: After oxidation of an aliquot of each sample with nitric acid, total iron was determined spectrophotometrically by the method of Yoe and Jones (5) using the red complex of Tirón and iron(III). The iron(III) in another aliquot was determined by reduction with an excess of potassium iodide and titration of the resulting iodine w'ith a standard solution of sodium thiosulfate.
Procedure. Weigh out a 300.0-mg. sample and dissolve it in distilled water slightly acidified with sulfuric acid. Dilute to 250 ml. in a volumetric flask. Each sample then should be analyzed immediately without interruption. Withdraw three 1-ml. aliquots and place each in a separate 25-ml. volumetric flask. Add 10 ml. of 0.3% 1,10-phenanthroline solution, buffer with 5 ml. of 0.2M potassium biphthalate solution, and dilute to the mark with distilled water. Read the absorbance of each solution at 396 µ and 512 µ as soon as possible and not later than 30 minutes after the complexes are
formed. Calculation of Results. Determine the concentration of total iron and the approximate concentration of iron(II) from standard concentration curves at 396 and 512 µ, respectively. Obtain the approximate concentration of iron(III) by difference. Find the absorbance value corresponding to this approximate concentration from the standard curve for iron(III) at 512 µ. Subtract this value from the observed absorbance at 512 µ to obtain
Table III.
Analyses of Prepared Samples Found®
_
Parts per million6 of Sample _% ConDiff. Method Method Method Method Diff. Sample A No. stituent Bd B Ac Bd B A Ac 0.04 13.67 0.07 1 6.56 6.60 13.74 Fe(total) -0.02 12.42 12.37 -0.05 5.96 5.94 Fe(II) 0.60 0.06 1.25 1.37 0.12 0.66 Fe(III) 0.27 2 5.89 6.02 0.13 12.27 12.54 Fe(total) -0.03 2.69 2.62 -0.07 1.29 1.26 Fe(II) 0.16 4.60 9.58 9.92 0.34 4.76 Fe(III) -0.27 -0.13 13.77 13.50 6.61 6.48 3 Fe(total) -0.17 7.98 7.63 -0.35 3.83 3.66 Fe(II) 0.08 0.04 5.79 5.87 2.78 2.82 Fe(III) 0.08 0.04 12.67 12.75 4 6.08 6.12 Fe(total) 3.46 -0.06 -0.03 3.52 1.69 1.66 Fe(II) 0.14 9.15 9.29 0.07 4.39 4.46 Fe(III) 0.10 0.05 14.48 14.58 6.95 7.00 5 Fe(total) -0.03 13.58 -0.01 13.61 6.53 6.52 Fe(II) 1.00 0.13 0.06 0.87 0.48 0.42 Fe(III) ® Average values from 3 aliquots of solution of sample. b Parts per million in 25 ml. of solution containing 1.2 mg. of sample, c Method A. Total iron and iron(II) determined with 1,10-phenanthroline. Iron(III) determined by difference. d Method B. Total iron determined with Tirón and iron(III) by iodidethiosulfate procedure. Iron(II) determined by difference. -
-
VOLUME
NO.
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JANUARY
1955
the corrected concentration of iron(II) from the appropriate standard curve. For the correct concentration of iron(III), subtract the corrected concentration of iron(II) from the concentration of total iron already determined. The results are not changed appreciably by a second approximation.
In Table III, the analyses of 5 prepared samples determined by the simultaneous spectrophotometric method with 1,10-phenanthroline are compared with analyses obtained by determining total iron with Tirón and iron(III) with the iodide-thiosulfate procedure. In columns 3 and 4, results are expressed as parts per million in a 25-ml. solution containing a V250 aliquot of 300 mg. of the sample. In columns 6 and 7, the compositions of the samples are given in percentages. In method B, the values given for parts per million of iron(III) were calculated from data obtained from the iodide-thiosulfate procedure using 50-ml. aliquots of a 250 ml. solution containing 300 mg. of sample. Values for iron(III) and iron(II) were gotten by difference in methods A and B, respectively. Values for parts per million of total iron and iron(II) in method A and for total iron in method B were obtained directly from absorbance measurements. The precision of the 1,10-phenanthroline method was determined by measuring the absorbances of solutions prepared with 0.5-, 1.0-, and 1.5-ml. aliquots of a solution of each sample. Duplicate determinations were made for each dilution. The average precision was 1.5% for total iron and 2.3% for iron(II). The maximum deviations were 5.4 and 5.3%, respectively. Discussion. Table III illustrates that results of analyses by the 1.10-phenanthroline method are in good agreement with results obtained independently by a method involving accepted
29
procedures of high accuracy. Comparison of results from these methods is particularly advantageous because in one case iron(II) is determined directly and iron(III) gotten by difference, whereas in the other case the determination of iron(III) is direct and that of iron(II) is by difference. The concentrations of total iron determined with 1,10-phenanthroline are in close agreement with values obtained with Tirón. This indicates that, although the absorbance coefficient of the 1,10-phenanthroline complexes at 396 µ is relatively small, absorbance measurements at this wave length give satisfactory results for total iron. The method presented in this paper is to be recommended for its simplicity. Two simultaneous spectrophotometric measurements on the same solution are sufficient for an analysis. No preliminary steps such as reduction, oxidation, or extraction of the sample are necessary. LITERATURE
CITED
(1) Blau, F., Monatsh. Chem., 19, 647-89 (1898). (2) Fortune, XV. B., and Mellon, M. G., Ind. Eng. Chem., Anal. Ed., 10, 60-3 (1938). (3) Harvey. A. E.. Jr., and Manning. D. L,, J. Am. Chem. Soc.. 74,
4744-46 (1952).
(4) XXralden, G. H., Hammett, L. P., and Chapman, R. P., Ibid., 55,
2649-54 (1933).
(5) Yoe, J. H., and Jones. A. L., Ind. Eng. Chem., Anal.
Ed.. 16,
111-15 (1944). Received for review May 28, 1954 Accepted October 14, 1954. Abstracted from a thesis submitted by John A. Smart in partial fulfillment of the requirements for the master of science degree, University of Arkansas. University of Arkansas Journal Series No. 1146.
Automatic Recording pH Instrumentation J. B.
NEILANDS and M. D. CANNON
Department of Biochemistry, University of California, Berkeley, Calif.
Titration
at constant pH offers great potentialities for following enzymatic activity, but because of lack of suitable instrumentation, it has been employed only for studies on certain hydrolytic enzymes. The apparatus described permits fully automatic determina-
tion of both ionization constants and volume of titrating fluid added as a function of time at constant pH. in enzyme research is illustrated Use of the instrument with acetyl esterase and lactic dehydrogenase. Auto-
is particularly matic, recording pH instrumentation useful for studying metal chelation and for following the course
of many
other chemical reactions, both cata-
lyzed and uncatalyzed. probably the most fundamental and useful electronic* instrument to be found in the biochemical laboratory. It also finds wide application in bacteriological and chemical laboratories as well as in plant industrial processes. Apart from simple measurements of the hydrogen ion concentration of solutions, two research applications of the pH meter are of paramount importance: the determination of the neutral equivalent and the ionization constants (pK„ values) of unknowm compounds and the measurement of volume of titrating fluid that must be added over a certain time interval in order to maintain a certain fixed pH. XXrhen the first of these techniques is applied to an unknown compound, information is obtained on the minimum molecular weight, the purity, and the possible structural features of the substance. The second application provides the data necessary for a kinetic analysis of a vast number of reactions, of both the catalyzed and uncatalyzed variety.
THE
pH meter
Both types of operation may, of course, be carried out manually, and in fact virtually all of such work has been so performed. The use of automatic recording devices, such as those described in this paper, greatly reduces the required experimental time, increases the sensitivity of the methods, avoids the human error, and makes it possible to follow reactions that are too rapid for manual methods. The instrument described is based on principles used by Lingane (6) and Jacobsen and Léonis (K)· However, certain modifications have been made in the design of the titration cell and in the method of adding the titrating fluid. The utility of the instrument in specialized biochemical problems, such as the study of enzymes, has been demonstrated. VARIABLE
is
pH TITRATION
shows the titration cell in detail
as well Figure outline of the other components of the system. A Viso-hp. synchronous Bodine motor is used to drive the buret plunger. The slow shaft speed of this motor is 6 r.p.m. The gear box and clutch assembly are shown in Figure 2. The slow-speed shaft of the motor is attached via a Lovejoy coupling to the fast-speed transmission shaft. The latter runs through a hollow shaft which bears a 16-tooth and a 32-tooth gear. The two shafts are connected with a keyway and the outside shaft may be slid back and forth over the inside shaft in order to position either the 16- or 32-tooth gear. The slow-speed transmission shaft carries a 120-tooth and a 100-tooth gear. One end of this shaft protrudes through the wall of the gear box and is attached directly to the clutch, which consists of a set of half nuts held by spring tension to the threaded buret drive shaft. When the motor turns, the fast and slow transmission shafts and the half nuts of the clutch are rotated. The clutch then pushes forward the threaded buret drive shaft, which cannot rotate. XXTien the
Apparatus.
as an
1
