Coulometric Reduction of Oxides on Tin Plate ROBERT
P.
FRANKENTHAL, THOMAS J. BUTLER, and RAYMOND T. DAVIS, Jr.
Applied Research laboratory, United States Steel Corp., Monroeville, Pa.
b In the coulometric reduction of oxides on tin-plate surfaces, the quantity of oxides measured i s influenced b y the composition of the electrolyte, the oxygen dissolved in the electrolyte, and the current density. It is dependent on the pH and the buffering capacity of the electrolyte. All electrolytes of pH 3 and 4 and buffered electrolytes of pH 5 to 7 give similar results, but unbuffered electrolytes of pH 5 to 7 and all basic electrolytes give a lower value. These low results are attributed to incomplete reduction of the oxides in the presence of a large concentration of hydroxyl ions a t the electrode surface, the probable reason being the dissolution of an alkaline-soluble intermediate hydroxide. Oxygen dissolved in the electrolyte causes a positive error in the measured quantity of oxide proportional to the square root of the time of electrolysis, indicating that the reduction of oxygen is a diffusioncontrolled process.
T
oxide films on metals greatly affect the surface properties of the metal. Numerous chemical and optical methods have been developed for evaluating the quantity of oxides on metal surfaces (3, 5, 10, 11). The coulometric principle was adapted by Miley (7’) to the measurement of oxide films on metals. It was modified by Campbell and Thomas (2) to make it more sensitive and applicable to very thin films. Salt and Thomas (8) and Katz (6) studied the quantitative determination of oxide films on tin by this method. However, it is believed that their use of a neutral, unbuffered electrolyte (potassium chloride) resulted in a negative error. Britton and Bright (f) have reported the use of a pH 7.3 buffered potassium chloride electrolyte. The effect of the composition of the electrolyte, specifically with respect to its p H and buffering capacity, on the quantitative determination of oxide films on tin plate has been investigated. The effect of oxygen dissolved in the electrolyte and the effect of current density have also been studied. HIN
EXPERIMENTAL
The electrolytic reduction cell was a block of hard rubber, bored out to hold 20 ml. of electrolyte solution. A silversilver chloride reference electrode and a
salt bridge were immersed in the electrolyte through holes in the top of the block. For the reference electrode a coil of silver wire with a large surface area of 1.9 square cm. was used to avoid polarization of the electrode. The salt bridge led to the silver-wire auxiliary anode. The sample was clamped against another hole on the end of the block so as to ensure a constant, reproducible cathode area. In the esperiments designed to study the effect of stirring, a beaker fitted with a Teflon stopper served as the cell. The electrodes, nitrogen bubbler, sample, and stirrer were brought through holes in the stopper. A constant current from a Millen No. 90201 direct current power supply was passed between the auxiliary anode and the cathode. The current density can be varied from 0.010 to 0.40 ma. per square cm. The potential of the cathode was measured with respect to the reference electrode. The potential-time curves were recorded on a Leeds &- Northrup Speedomax Model S, Type G, recorder (10-mv. full scale; chart speed of 2 inches per minute) or on a Leeds & Northrup Speedomax high speed recorder Type G (10-mv. full scale; chart speed of 60 inches per minute). A voltage divider of 150 kohms was in parallel with the recorder to reduce the input voltage to less than 10 mv. The current density was measured with a Model S, Sensitive Research Instrument Corp. milliammeter, and was corrected for losses through the voltage divider. No correction was made for the I R drop through the cell because it was less than the experimental error in the potential measurements. All measurements were made with air-formed oxides on tin plate and were performed a t least in duplicate. The tin plate was vapor degreased in trichloroethylene before the oxides were reduced. Reagent grade chemicals were used without further purification, and all solutions were prepared with distilled water. The work with deaerated solutions was performed in an inert-atmosphere box through which high purity nitrogen (Seaford grade, Airco, Air Reduction Sales Co., New York) was passed. Oxygen could not be detected by mass spectrometric analysis in this nitrogen. The nitrogen was also bubbled through the solutions prior to their use. A Beckman Model H pH meter was used to determine the pH of all solutions. RESULTS
The potential-time curve observed is similar to the familiar S-shaped poten-
tiometric titration curve. The initial potential corresponded to that required for the reduction of the oxide. A shift to the hydrogen-evolution potential was observed when all the oxide had been reduced. The inflection point on the curve was taken as the end point of the determination. Because the composition of the tin oxides was unknown, the quantity of surface oxides (oxide-film value) was expressed in millicoulombs per square centimeter (current density X time). Effect of Electrolyte Composition. The oxide-film values determined in 0.1M potassium chloride electrolyte were always lower than those determined in 0.001M hydrobromic acid electrolyte. This difference was not due t o attack by the electrolytes, because tests showed that the oxides were not attacked by the electrolytes o w r periods of time much greater than those required to measure the oxide film. To determine whether the anion, cation, or differences in concentrations were responsible for the conflicting results in the two electrolytes, oxide-film values were measured in potassium chloride and potassium bromide solutions at two concentrations and in 0.001M hydrochloric acid and 0.001M hydrobromic acid. The results are summarized in Table I. The current density was 0.077 ma. per square em. Acids 0.1M in concentration could not be used because they attacked the oxides. The data indicate that the hydrogen ion concentration is of prime importance, whereas the nature of the anion and the concentration of the electrolyte have little observable effect. To determine the effect of pH on the oxide-film value, the following electrolytes were used: buffers in the range pH 3 to 9.6; unbuffered hydrobromic acid in the range p H 3 to 6; sodium hydroxide of pH 13 (0.1M); and potassium chloride of pH 7 (0.1M). Tests showed that no oside was dissolved by the electrolytes which were used during the time required to run the reduction. All the buffers (unless otherwise specified) were prepared from 0.1M citric acid and 0.2M disodium hydrogen phosphate. The hydrobromic acid solutions of pH 4, 5 , and 6 m-ere made 0.001M in potassium chloride to increase their conductivity. The results are plotted in Figure 1. When some phenolphthalein was added to the potassium VOL. 30, NO. 3, M A R C H 1958
441
electrolytes, the excess hydroxyl ions are neutralized, but not in a slightly acidic unbuffered medium where the conditions a t the electrode surface will simulate those in a basic solution. This is supported by the experiment with phenolphthalein in the potassium chlcride electrolyte. Because tin oxides are amphoteric, low results in the basic and in the unbuffered slightly acidic electrolytes can be explained by the dissolution of some of the oxides in this alkaline medium or by the dissolution of a hydroxide intermediate in the reduction process. Stirring should reduce the alkalinity a t the electrode surface and should prevent the dissolution of the oxide. This !vas observed with the hydrobromic acid, p H 5. Not only did the oxide-film value increase to that value measured in the acidic and in the buffered electrolytes, but the oxide and the hydrogen reduction potentials approached values close to those calculated from the Nernst equation (E = E, -RT/nF A pH, where E , was taken as the potential measured in hydrobromic acid of pH 3). Stirring in the potassium chloride did not cause an increase in the oxide-film value, and the positive shift in the reduction potentials was not observed because the stirring w.as not sufficiently effective in an electrolyte with a pH on the border between those that give only the lower oxide-film value and those that can give the higher one (see Figure 1). Khen very violent agitation was applied, hoivever, the reduction potentials did shift in the positive direction. The intermediate value observed with citratephosphate buffer of pH 7 (Figure 1) results from a poor buffering capacity a t the borderline pH. According to Stern (9) the pH a t the electrode surface may reach a value bc-
Table I.
Oxide-Film Values as Influenced b y Composition of Electrolyte" 01.1.1' KC1 0.1M KBr 0.001.W KC1 0.001M KBr O.0OlA1I'HCl 0 . OOlM HBr 1 . 6 h O . 1 1.6 k O . 1 1.6 5 0 . 1 1.82 z O . l 2.6 f O . l 2.6 5 0 . 1 a Rfilliconlombs per square cm.
Table II. Oxide-Reduction Potentials and Hydrogen Evolution Potentials"
Electrolyte HBr, pH 3 HBr, pH 3 HBr, p H 5 HBr, pH 5 KC1, pH 7 KCl, pH 7 NaOH,pH13 KaOH,pH 13
r'olts trode. a
us.
Stirring
KO
Yes
NO
Yes NO
Yes No
Eo -0.60 -0.60 -1.03
-0.71 -1.04 -1.02 -1.18
EH -0.95 -0.95
-1.37 -1.06 -1.40 -1.38 -1.54 -1.54
Yes -1.18 silver-silver chloride elec-
chloride, a red color appeared a t the electrode surface during the electrolysis, indicating a pH greater than 9. Effect of Stirring. The experiments were repeated with stirring of the electrolyte in hydrobromic acid (pH 3 and 5 ) , potassium chloride, and sodium hydroxide solutions. I n hydrobromic acid of pH 5 the oxide-film value increased t o that observed in hydrobromic acid of pH 3, while in the other media they were unaffected by stirring. Potential Measurements. The oxide-reduction potentials, E,, and the hydrogen-evolution potentials, Ex, in the different electrolytes are tabulated in Table 11. The current density was 0.100 ma. per square cm. The potentials are a function of the current density and cannot be related to thermodynamic potentials. The variation of the potentials with current density in the various media is uniform. In another experiment, the hydrogenevolution potential tw-s measured in the potassium chloride electrolyte as a function of the degree of stirring. When very violent agitation was appliede.g., the sample was rotated, a magnetic stirrer n.as used, and nitrogen was bubbled through the cell-a potential of approximately -0.9 volt mas observed. The oxide-film value could not be measured under these conditions because the nitrogen bubbles would adhere to the sample, resulting in incomplete reduction. When the stirring mas discontinued, the potential returned to the more negative value given in Table 11. When a few drops of concentrated hydrochloric acid were added, raising the acidity to a pH of 2.1, the hydrogenevolution potential immediately dropped to -0.90 w l t and mas not affected by stirring. DISCUSSION
Figure 1 shows that oxide-film values measured in all acidic electrolytes of pH 442
ANALYTICAL CHEMISTRY
4 or less or in buffered electrolytes of pH 7 or less were the same within experimental error. The values determined in basic media or in unbuffered electrolytes in the range of pH 5 to 7 agreed among themselves but were lower than those observed in the acidic and buffered electrolytes. The break point between the low and high results occurred between pH 7 and pH 8 in the buffered electrolytes and between pH 4 and pH 5 in unbuffered electrolytes. Although an independent measure of the oxide-film value is not available and it is not certain whether the high or the low value, if either, is correct, certain conclusions indicating that the higher value is the more accurate one can be deduced from the available data. If the results in the highly acidic and in the buffered electrolytes were high and the results in the other media were correct, an electrolytic reaction that could account for the high rcsults would have t o exist. The only reductions that could take place (other than those of the oxides) are the evolution of hydrogen and the reduction of tin t o stannane. Both these reactions should be p H dependent. Because no such dependency mas established, it is highly improbable that the results in the acidic electrolytes are high. Low results in the more alkaline electrolytes can be explained simply. During the electrolytic reduction of an oxide, hydrogen ions are consumed, thus increasing the hydroxyl ion concentration at the elwtrode surface. In highly acidic or in slightly acidic buffered
= t
A
I 0
+
o
A
4 -I
+
HBr
A CITRATE- PHOSPHATE 0 BUFFERS KCI
:o,*[
s
NaOH
v COMMERCIAL
I
2 ~
A
PHOSPHATE BUFFER I
I
I
I
I
I
I
tween 12 and 14 during a reduction process. I n the present reduction potential data, a value between p H 10 and p H 11 is estimated. inasmuch as the reduction potentials measured in the 0.1M sodium hydroxide are 140 mv. more negative than those measured in the unstirred potassiuni chloride. Effect of Dissolved Oxygen and Current Density. The ouide-film value was measured in deaerated and air-saturated solutions, as a function of current density (0.019, 0.039, 0.078. 0.116, 0,155, 0.233, 0.310, and 0.38s ma. per square cm.) and on various tin-plate specimens with different quantities of surface oxides. I n deaerated solutions, the plot of oxide-film value us. current density approached a straight line parallel to the abscissa. Slight deviations caused by a trace of oxygen remaining in the solution were observed. The error in the oxide-film value due t o the presence of oxygen in the electrolyte in the air-saturated solutions increased nith decreasing current density for a particular oxide level. Similarly, the error increased with increasing oxidefilm value when the current density was held constant. This is consistent with the assumption t h a t oxygen was reduced concurrently with the oxide. For long electrolysis times-Le., for low current densities or high oxide-film values-the quantity of oxygen reduced was greater than that reduced during short times of electrolysis, hence, the greater error. This indicated that the reduction of the oxygen could br a diffusion-controlled process
(at least for the times of electrolysis measured in these experiments). Assuming linear diffusion, the current, i, a t any time, t, is given by (4) it = nFK-l~2ADl/2Ct--112 (1) where A is the electrode area, D is the diffusion coefficient of the oxygen, C is the bulk concentration of the oxygen, and the other terms have their usual significance. To determine the quantity of electricity that flows over a period of time, Equation l must be integrated over the time interval. Thus Q =
2nF=-1/2AD1/2Ctl
2
(2)
where Q is the quantity of electricity, which is a measure of the quantity of oxygen reduced. For a diffusion-controlled process, a plot of the error due to the presence of dissolved oxygen in the electrolyte us. the square root of the time of electrolysis should give a straight line. A plot of the data gives a straight line, in agreement with the theory. Stirring the electrolyte also increased the error; this is consistent with a diffusion-controlled process. At sufficiently high current densities, the oxygen error became almost negligible, especially for tin plate that only had a very thin oxide film. SUMMARY
The quantitative coulometric determination of tin oxides is affected by the p H and the buffering capacity of the electrolyte. Results in unbuffered electrolytes of pH 5 or greater and in all
electrolytes of pH 8 or greater are low because of the dissolution of some of the oxides, the dissolution being caused by the high alkalinity at the electrode surface. The p H a t the electrode surface, under these conditions, is estimated to he between 10 and 11. Oxygen dissolved in the electrolyte causes an error proportional to the square root of the time of electrolysis. The evidence indicates that the rate of reduction of the oxygen is diffusion controlled. LITERATURE CITED
(1) Britton, S. C., Bright, K., Xetallurgia 56, 163 (1957). (2) Campbell, FV. E., Thomas, U. B., Trans. Electrochem. Soc. 76. 303 (1939). (3) Constable, F. H., Proc. Roy. SOC. (London) A117, 376, 386 (1928). (4) Delahay, P., “Kern7 Instrumental
Methods
in Electrochemistry,”
p. 51, Interscience, Xew Pork, 1954.
(5) Evans, U. R., Stockdale, J., J . Chem. SOC.1929, 2651. (6) Katz, W., Stahl IC. Eisen 76, 1672 (1956). ( 7 ) Miley, H. .4.,Iron Steel Inst. Carnegie Schol. Mem. 25, 197 (1936). (8) Salt, F. K., Thomas, J. G. N., Sature 178. 434 (1956). (9) Stern,’ M., J. El&trochern. SOC.102, 609 (1955). (10) Tammann, G., Bochow, K., 2. anorg. Chem. 169, 42 (1928). (11) Vernon, R H. J., Trans. Faraday SOC.31, 1668 (1935).
RECEIVEDfor review May 29, 1957. Accepted October 5, 1957. Division of Physical and Inorganic Chemistry, 130th Meeting, ACS, Atlantic City, K. J., September 1956.
Infrared Identification of Some Sulfur Derivatives of Long-chain Fatty Acids HEINO SUSI,
N.H. KOENIG, W . E.
PARKER, and DANIEL SWERN
Eastern Regional Research laboratory, U.
b Infrared absorption spectra of some long-chain fatty acid derivatives containing sulfide, sulfoxide, and sulfone groups were studied. The spectra were investigated to determine the effects of such substitutions on the characteristic bands of fatty acids and to establish means for detection and determination of these compounds in the presence of each other. Interruption of the carbon choin b y -S-, -SO-, or -SO*makes the band progression region less useful for chain-length determination, but introduces spectral detail that can be used to identify individual compounds. Very closely related sulfur derivatives of fatty acids can be
S. Department of Agriculture,
Philadelphia 7 8, Pa.
distinguished b y studying the 1350to 1 180-cm.-’ region of solid-state spectra obtained from potassium bromide pellets. For differentiating between the main classes, it i s best to utilize the extremely regular dilute solution spectra of the corresponding methyl esters.
T
HE: ASALYTICAL USEFULKESS of the infrared absorption spectra of longchain fatty acids has been discussed and evaluated by a number of workers, most recently by Jones, LIcKay, and Sinclair (4, 12) and by Aleiklejohn, Meyer, Aronovic, Schuette, and Xeloche ( 7 ) . The spectra show bands that are characteristic for the carboxyl and
methylene groups and for the methyl end groups. In the crystalline state the “band progression,” a regular series of evenly spaced absorption maxima, is observed in the 1350- to 1180-cm.-l region. This band progression frequently provides the only practical means of spectrally differentiating various saturated fatty acids because their spectra are otherwise similar. The present study describes the spectral effects produced by introducing a sulfide, sulfoxide, or sulfone group into the fatty acid chain. Some of the characteristic frequencies of these groups have been established ( 1 , 9, IO). It can hardly be expected, however, that introduction of bulky or highly polar VOL. 30, NO. 3, MARCH 1958
443
