Trigger Effects in Spontaneous Electrolysis

TRIGGER EFFECTS IN SPONTANEOUS ELECTROLYSIS

DISSERTATION

Presented in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy In the Graduate School of the Ohio State university

By

RICHARD MAC WILSON, B.S., M.S.

The Ohio State University 1959

Approved by

Adviser Department of Chemistry Dedicated to m y wife and family

11 PREFACE

The research described herein concerning the galvanic cells established between two similar electrodes with the same history suggests a problem In corrosion not previously considered, that is, electrochemical corrosion between two like electrodes. The magnitude of this corrosion is relatively small but It would, under proper conditions, be sufficient to initiate cavitation and lead to serious damage of the corroding object.

Ill ACKNOWLEDGMENT

The author wishes to express his sincere appreci­ ation to Dr. William M, MacNevin for his encouragement and guidance throughout the course of this research.

iv CONTENTS Page INTRODUCTION ...... 1 Electrochemical Cells ...... 1 Concentration Cells* ...... 3 Thermogalvanlc Cells ...... 6 Gravity Cells ...... 8 Centrifugal Cells ...... 13 Crystallographlc Effect on Potential . . . 15 Stress Corrosion ...... IT Motor Electric Potential ...... 23 Effect of a Magnetic Field ...... 25 Effect of Interfaclal Free Energy .... 27 Photovoltaic Cells ...... 28 Properties of Aluminum ...... 30 History of the Aluminum-Platinum Electrode Pair 40 Properties of Zirconium ...... 50 STATEMENT OF PROBLEM ...... 54 EXPERIMENTAL APPARATUS...... 56 EXPERIMENTAL PROCEDURE ...... 65 RESULTS ...... 69 Aluminum-Platinum Electrode P a i r ...... 69 Aluminum-Aluminum Electrode P a i r ...... 82 Zirconium-Platinum Electrode P a i r ...... 148 Zirconium-Zirconium Electrode P a i r ...... 148 DISCUSSION...... 154 SUMMARY ...... 168 BIBLIOGRAPHY...... 172 AUTOBIOGRAPHY ...... 178 v LIST OP TABLES Table Page 1. Effect of centrifugal force on the potential of the iodine electrode...... 15 2• Motor electric potential of the copper e l e c t r o d e ...... 24 3. Photovoltaic effect on the copper-copper oxide e l e c t r o d e ...... 29 4. Properties of aluminum...... 30 5. Partial list of the electromotive series . . . 31 6 . Comparison of the heat of formation of some sodium, aluminum and iron compounds...... 32 7. Corrosion resistance of zirconium ...... 51 8 . Analysis of reagent-grade aluminum wire . . . 61 9. Analysis of zirconium rod...... 62 10. Effect of the size relationship on the response of the aluminum-aluminum electrode pair to 20 micrograms of sodium fluoride...... 94 11. Effect of pH on the response of the aluminum- aluminum electrode pair to 20 micrograms of sodium fluoride...... 95 12. Response of the same set of 2:1 aluminum electrodes to additions of 20 micrograms of sodium fluoride...... 105 13. Reproducing conditions on an aluminura-platlnum electrode pair and the response to 20 micro­ grams of fluoride. Pour spirals (4.8 square centimeters) of aluminum wire...... 117 14. Reproducing conditions on an aluminum-platinum electrode pair and the response to 20 micro­ grams of fluoride. A continuation of Runs No. 268-271...... 118

vi vii LIST OF TABLES (Contd.) Table Page 15* Response of the aluminum-calomel electrode pair to 20 micrograms of fluoride. All solutions were outgassed with nitrogen and a nitrogen atmosphere maintained above the cell solution...... 127 16. Effect of outgassing the cell solution on the response of the aluminum electrode to 20 micrograms of fluoride ...... 128 17. Effect of position on the behavior of the 2 :1 aluminum-aluminum electrodepa i r ...... 129 18. Response to 20 micrograms of fluoride of selected combinations of four similar aluminum electrodes...... 134 19. Effect of shape on the behavior of two aluminum electrodes with relative surface areas of 1 .8 and 2 .8 square centimeters ...... 135 LIST OP FIGURES Figure Page 1. Current-temperature relationship in the aluminum thermogalvanlc c e l l ...... 10 2. Baker type curve ...... 43 3. Holland type curve ...... 43 4. Actual size diagram of Cell I and an electrode holder used with this cell ...... 59 5. Comparison of the response to 20 micrograms of sodium fluoride of the Baker electrode at 27° C., a spiral electrode (99# Al) at 25° C. and a spiral electrode (99# Al) at 50° C ...... 72 6 . Log 1 versus t for the current decay resulting from the addition of 20 micrograms of sodium fluoride to 3 5 .0 mis of 0 .2 M acetic acid at 50.0° C. with the aluminum-platinum electrode p a i r ...... 74 7. Fluoride calibration curve with the aluminum- platinum (99.99# Al) electrode pair...... 76 8 . Effect of the preliminary treatment, 3 minutes with 0.01 M HF, on the aluminum anode In combination with a platinum c a t h o d e ...... 79 9. Effect of a soluble aluminum salt on the galvanic current flow between the aluminum- platinum electrode in the presence of 50 micrograms of sodium fluoride...... 8 l 10. Sample current-time curve obtained with a large aluminum electrode in contact with a platinum cathode In 0.2 M acetic acid...... 85 11. Sample current-time curve obtained with a small aluminum electrode in contact with a platinum cathode in 0 .2 M acetic acid...... 85 12. Current-time curve for a 10:1 ratio of an aluminum-aluminum electrode pair in the presence of 20 micrograms of sodium fluoride . 87 viii lx LIST OP FIGURES (Contd.) Figure Page 13. Current-time curve with the current values obtained with the Standard Research Instru­ ments Company milliammeter. 20 micrograms of sodium fluoride added...... 91 llf. Effect of successive additions of fluoride on the current flow between a 2 :1 aluminum- aluminum electrode p a i r ...... 93 15. Effect of pH on the response of the alumi- num-aluminum electrode pair to added fluoride 98 16. Effect of adding fluoride to the center of the cell solution, to the vicinity of the anode and to the vicinity of the cathode. . . 101 17. Variation of the response of a 2:1 aluminum- aluminum electrode pair to changes In added f l u o r i d e ...... 104 18. Effect of use on the response of the aluminum-aluminum electrode pair to added f l u o r i d e ...... 107 19. Current response of the aluminum-aluminum electrode pair to a constant fluoride concentration ...... 114 20. Current response of the aluminum-platinum electrode pair to a constant fluoride concentration ...... 116 21. Current flow between a platinum end saturated calomel electrode in 0.2 M acetic acid. The platinum was either rinsed with water or Immersed In hot concentrated nitric acid. • . 122 22. Reproducible current flow between a platinum and saturated calomel electrode in 0.2 M acetic acid. The platinum was stored in the electrolyte before use and the solution out- gassed...... 125 X LIST OP FIGURES (Contd.) Figure Page 2 3. Effect of differential stirring on the current flow between two similar aluminum electrodes. . 132 24. Change of potential with time of a large and small aluminum electrode In 0.2 M acetic acid • 137 25. Diagrams of the spiral and disk-shaped elec­ trodes and the apparatus used for measurements with the Luggin capillary...... 141 26. Change of resistance with time of a 0.2 M acetic acid solution containing 2 .0 grams of aluminum metal •••••...... 145 27. Effect of ultrasonics on the current flow between two aluminum electrodes In 0.2 M acetic acid...... 147 2 8 . Fluoride calibration curve with the zirconlum- platinum electrode pair In 0.1 M perchloric acid...... 150 29. Current-time curve for two zirconium electrodes in 0.1 M perchloric acid...... 153 30. Polarization diagram for a process under anodic c o n t r o l ...... 162 INTRODUCTION

Electrochemical Cells

The tern electrolysis means in general terms "a chemical change due to the passage of electricity." A common example of this Is the battery-operated electroly­ sis cell In which chemical reactions are initiated and maintained by the passage of current from an external current source. The electricity referred to is direct current or Ohm*B Law electricity, and Faraday!s Laws re­ late the quantity of electricity flowing through the cell to the amount of chemical change. In contrast, cells are known that do not require an outside current source. These cells are known by several names: galvanic cells, voltaic cells, spontaneous electrolysis cells and Internal electrolysis cells. In such a cell the current flow does not cause the chemical reactions to occur but Is the result of chemical energy or a related form of potential energy, generated by a spontaneous chemical reaction, being converted Into electri­ cal energy. A common example of this is the Daniel Cell in which electrodes of metallic zinc and copper are Immersed in solutions of zinc and copper sulfates respectively. If 1 these two metals are connected externally by a metallic conductor, a current will flow through the conductor. The half cell reactions are

Zn° --- > Zri4-*' + 2e“ (anode) Cu*"*" + 2e“ — ■■ > Cu° . (cathode)

In this cell the electrons flow from the zinc to the copper and if the activities of the zinc and the copper Ions in the solution were each unity, a potential difference of 1 ,1 0 5 8 volts would exist. Cells such as this are quite useful in electro­ chemical analysis since they furnish, in addition to a current source, a method of controlling anode and cathode potentials. For example, Ullgren, 1 as early as 1868, demonstrated that copper could be determined electro- chemlcally without an external current source by using a cell consisting of the zinc and copper electrodes.

(1) C. Ullgren, Z. Anal, Chem, Zj **2 (1868).

In any electrochemical cell the reaction at one electrode Is reduction (gam of electrons) and at the other electrode oxidation (loss of electrons). The term anode Is used to designate the electrode at which oxidation takes place, cathode the electrode at which reduction takes place. This designation of electrodes is correct for any 3 electrochemical cell, although there Is a difference in the relative signs (+ or -) of the electrodes. In a gal­ vanic cell the anode is negative and the cathode Is posi­ tive. In a battery operated cell the signs are reversed, that Is, the anode Is positive and the cathode Is negative.

Concentration Cells

The previous example of a galvanic cell was one consisting of two unlike electrodes in which the cell processes were chemical reactions. This, however, need not be a necessary requirement of a galvanic cell; for example, there are cells known that are composed of two like electrodes and the net cell process Is not a chemical reaction but rather a physical transfer of a substance from one phase to another. Such cells are called "concen­ tration cells" and the potential Is a function of the concentrations In the two phases being transferred. A typical example of a concentration cell Is the cell consisting of two dilute amalgam electrodes of different concentrations, and thus different activities, In a solution of a salt.

M (Hg) | rf* I M (He) C, C- ' C„ 4 If In thlB cell the concentration of metal In C^ Is greater than in Cg, the net process Is merely a transfer of metal from concentration to concentration Cg. A similar type of cell may be made up with electrodes of a pure metal and an amalgam. Concentration cells ere laiown that are dependent on a difference in pressure in which the net process is the transfer of a gas from one cell to another.

H2 (1 atm) | ft* | Hg (2 atm)

In addition to the two previous types of concen­ tration cells there Is another type that Is dependent on the difference In concentration of the reacting Ionic species in solution.

Ag | Ag+ NO^ (C^) | Ag+ NO^ (C2) | Ag

In this cell the net process Is the transfer of Ions from the concentrated to the dilute solution. The same effect can be produced In a single cell by prepolarlzlng two similar electrodes with an external current source. This causes a change m the concentration of the reacting Ions in the Immediate neighborhood of the two electrodes. When this external current source is re­ moved and the electrodes shorted a current will flow be­ cause of the unequal concentration of the Ions in the 5 Immediate vicinity of the electrodes. This is actually a concentration cell and the current flow decays to zero in a short period of time because of the restoration of equilibrium conditions in the cell. A pseudo-concentration cell is set up between two cells containing the same concentration of reacting species but a difference in the total ionic strength of the two cells. In such a cell the effective concentration, that is activity, would be different because of a differ­ ence in the activity coefficients (a = y c) and the net effect would be a transfer of ions to the cell of lower activity. Concentration cells are widely used in potentio- metric techniques; for example, measurement of pH with a pH meter involves a concentration cell dependent on a difference in hydrogen ion concentrations. The fact that at equilibrium a concentration cell has no potential difference has been used to determine transition points. Tn a cell made up of two allotroplc forms of a metal in a solution of one of its salts, a potential difference exists provided the temperature is not at the transition point. Only at this temperature are the two forms in equilibrium and the potential differ­ ence zero. 6 Thermogalvanlc Cells

A thermogalvanlc cell is a form of a galvanic cell In which a metal Is removed from one surface and deposited on another as a result of a temperature difference between two like electrodes* In such a cell It seems reasonable to assume that the only important source of electrical work Is the reversible transfer of latent heat (TAS), of the simple electrode reaction, from a higher to a lower temperature• For such a cell with electrode temperatures differing by dT degrees Kelvin,

w = nFdE = TA SdT/T = A SdT or dE/dT = A s/nF.

The reversible potential Er of any cell may be estimated by integrating this equation between the two temperatures of the electrodes* When this gives a positive value for Er the electrode at the higher tem­ perature will be the anode and vice versa. It has been reported that at a temperature differ­ ence of lf>0° F. a potential difference of 80.5 millivolts 2 Is developed between two copper electrodes. The warmer electrode Is the cathode.

(2) N. E. Berry, Corrosion 2, 261-7 (19^6). 7 Similar results have been noted with lead and predicted for zinc, Iron, cadmium and tin. 2 Similar behavior has also been reported for alumi­ num although aluminum cannot be deposited from an aqueous •2-c; solution ,J **

(3) L. Kahlenberg and S. J. French, Trans, Am, Electrochem. Soc. 52. 355 (1927). (4) R. B, Mears and R. H, Brown, Ind. Eng, Chem, 3 3, 1001 (1941). (5) C. Crussard and F. Aubertln, Rev, Met. 45. 402 (1948).

Pltzer^ reported that two aluminum cans (94.4# Al) of about 45 square centimeters surface area developed a galvanic current of the order of only a few microamperes until the temperature difference reached 60° G. Above

(6 ) E. C. Pltzer, J. Electrochem. Soc. 104. 70-74 (1957).

90° C. there was a marked Increase In both the potential and galvanic current, m the aluminum thermogalvanlc cell the warmer electrode Is anodic. One of the complicating factors encountered in studying a metal Buch as aluminum Is the oxide film. Alumi­ num oxide Is more soluble in the warmer solution. This solubility difference could play a very significant if not dominant role in the determination of the polarity of the cell, that la, the Increased solubility In the warmer solution would decrease the thickness of the oxide layer and therefore decrease the passivity of the warmer elec­ trode, thus making that electrode more susceptible to attack. The large increase m potential above 90° C. is probably due to the transition of fi AlgO^'S^O to «c AlgO-j*31*20 which is known to take place In this temper­ ature range. Figure 1 shows a current-time curve obtained by Pitzer. 7 Nickel thermogalvanlc cells have been reported. The warmer electrode is anodic.

(7) D. S. Carr and C. F. Bonilla, J. Electrochem. Soc. 22, 475 (1952).

Since nickel Is also known to be coated with an oxide layer, this polarity behavior can probably be ex­ plained in terms of the difference in solubility of the oxide layer In cold and warm solutions.

Gravity Cells

The effect of the force of gravity on the potential of a galvanic cell Is usually negligible. However, there are examples known In which the effect Is appreciable. 9

Pig. 1.— Current-temperature relationship in the aluminum thermogalvanlc cell. CURRENT i\ua) 900 800 500 400 IOO 0 0 6 0 0 2 700 300 40 - - - 080 50 EPRTR DFEEC (°C) DIFFERENCE TEMPERATURE 90 too 10 11 Experiments undertaken to study the effect of difference in a gravitational field on the potential of 8*9 a galvanic cell have teen carried out by Des Coudres 10 and Tolman.

(8) T. Des Coudres, Ann. der Physik u. Chemle 49. 284 (1892).

232 (189 ^?! * * ^ (10) R. C. Tolman, Proc. Am. Acad. Arts Sci. 46. 109 (1910).

Two reversible calomel cells were connected by a rubber tube filled with chloride solution so that the relative heights of the two electrodes could be altered.^

Hg | HgCl, MCI | MCI, HgCl | Hg A h = 0 h = h B

(Note- at the time of this work calomel was believed to be HgCl.) In such a cell the following are applicable:

- 4 Z : EF 5 tn AUnCl - A U Ci(8)- Z ® Gibbs free energy

E - electromotive force

P “ Faraday

tjn ~ transference number 12

ujjci * chemical potential of the solute

uCl(s) “ chemical potential of the combined chloride

Using the relationship du^ = (M^V^pJg dh, the above equation becomes

EF * gh[ tm (M[(cl -VMCl/»)-(Mci -Vci(o)/0 )] .10 g = gravitational force per gram

h = height

M s molecular weight of the substance

/° * density of the solution

V = partial molal volume

For the passage of one Faraday through this cell the following changes would take place in the cell:

1 , the transfer of tm number of equivalents of the salt, MCI, from the region of the lower to the higher electrode 2. the reaction of mercury with an equivalent of chloride Ion at the lower electrode (oxidation) 3* the reduction of an equivalent amount of HgCl at the higher electrode That Is, the transfer of an equivalent of com­ bined chloride from the higher to the lower cell* 13 The following is a partial summary of experimental results obtained in gravity cells such as those described above.^

Salt Molality E.M.F.. volts/cm x 10^ BaCl2 O .9 8 1.70 KC1 2,71 0.51

Centrifugal Cells

The effect described in the gravity cell Is greatly increased if a centrifugal field is substituted for the gravitational field.

(11) R. C. Tolman, J. Am. Chem. Soc. 32, 121 (1911)* (12) F. 0. Koenig and S. W. Orinnell, J. Phys. Chem. 44. 463 (1940). (13) S. W. Qrinnell and F. 0. Koenig, J. Am. Chem. Soc. £4> 682-6 (1942). (14) D. A. Maclnnis and B. R. Ray, J. Am. Chem. Soc. II, 2987-9 (1949). (15) D. A. Maclnnis, Proc. Am. Phil. Soc. 97, 51-5 (1953).

MacinniB 0 gives a summary of the early work in this field.

(16) D. A. Maclnnis, "The Principles of Electro­ chemistry," Reinhold Publishing Corporation, New York, N.Y., 193a PP-178-180. 14 The effect of the centrifugal force may be shown as follows:

n - revolutions per second

r ^ r g “ radii In the centrifugal field

Tolman and Maclnnis used the cell

(Pt) | I2 , KI, I2 | (Pt)

Using the same approach that was applied to gravity cells the following equation was obtained: 11 EP = Z 1 ? n 2 (r? - rf) [tK (MKI - VKX/o) - (»% - V ^ ) ] .

E = electromotive force

P - Paraday

tj^ - transference number of the cation

/° = density of the solution

mKI = molecular weight of potassium iodide

Mj * atomic weight of iodine

VKT - partial raolal volume of potassium iodide

Vj z partial atomic volume of iodine 15 The following experimental results were obtained for the cell

11 (Pt) | I2, Lil, I2 | (Pt):

r = 29.4 cm* r - 4.2 cm.

Table 1.-Effect of centrifugal force on the potential of the Iodine electrode.

n E x 10-3 volts 52.3 3.23 56.2 3.72 59.2 4.16 6 3 .8 4 .8 0 68.3 5.51 72.4 6.22

Measurements with gravity and centrifugal cells may be used to calculate transference numbers. Comparison of the values obtained in this manner with proven methods shows that the results are very nearly correct.

Crystallographlc Effect on Potential

The electrochemical anisotropy of crystalline solids has been established by several authors. Tragert and Robertson*^ give a historical review of the points leading up to this conclusion, included In this are rates of

(17) W. E. Tragert and W. D. Robertson, J. Electrochem. Soc. 102. 86 (1955). 16 solution of Ionic crystals and chemical reactivity of single crystals as a function of the lattice orientation. The experimental evidence indicates that the crystallo- graphlc dependence of electrochemical phenomena is real. It was because of this that some investigators chose to work with two-phase copper amalgam electrodes to obtain reproducible results. In the light of this It is clearly evident that this effect should be taken Into consideration when determining standard electrode potentials. The following Is a partial summary of experi­ mental results obtained by Tragert and Robertson and a list of the conclusions drawn from this work:

M (CuSO)j) Temp* (°C.) Crystal Plane (Cu) E (v )

0.01447 2 5 .0 111 0.40890 110 0.41273 100 0.41262 210 0.41236

1. The oxidation potential of solid copper, and presumably of other metals, Is dependent on the crystallo- graphlc configuration of the atoms on the surface* 2. The only Btable plane of the face centered cubic system is the close packed octahedral plane and accordingly only this plane produces a reversible cell. This electrode reaction may thus be written ^(8)111--- > Cu++ + 2e~. 17 3* All other crystal plane surfaces are metastable and approach the (ill) configuration in time, 4. Polycrystalline electrodes approach the stable state providing enough time Is allowed for the establish­ ment of the stable crystallographlc configuration.

Considering this information it is evident that there definitely would be a potential difference between two unlike crystal faces and thus the probability of a galvanic cell between two anisotropic crystalline forms if this situation could be realized experimentally. Morize^ reported that A1 (ill) is 50 millivolts more negative than A1 (001).

A1 (111) = - 1.20 volts A1 (001) = - 1.15 volts

(18) P. Morize, Metaux et Corrosion 22. 71 (1947).

Stress Corrosion

Stress corrosion can be defined in a general way as the acceleration of corrosion by stress.*9

(19) M. 0. Fontana, md. Eng. Chem. M(No. 3), 99-100A (1953). 18 There are several different effects that alter the corrosion rate of a metal under stress, the most sig­ nificant of these are the cracking of a protective coating on the metal, plastic deformation, and elastic deformation. Plastic deformation involves a stress great enough to ex­ ceed the elastic limit of the metal and thus cause slippage or rupture of the crystal structure within the metal. Therefore, permanent damage Is done to the metal and initial conditions are not restored on removal of the stress. Elastic deformation Involves a stress within the elastic limit of the metal so that on removal of the stress the metal returns to the original state. There are several theories concerning the mecha­ nism of stress corrosion:

The electrochemical theory The mechanical theory The strain accelerated decomposition theory The anodic shift theory The film theory

These theories are all described by Fontana; however, the question of Just which theory gives the best representation of the mechanism has yet to be settled. The current belief Is that a combination of these theories is needed to explain all cases of stress corrosion. 19 Of these causative factors the first one, that is, rupture of a protective coating, seems to be the most significant, at least as far as the magnitude of the stress 20-25 voltage is concerned* This effect has been widely studied*

(20) E. M* Zaretskii, Doklady Akad* Nauk. SSSR 53,77-9 (19*7). (21) H. L* Logan, J. Res* Nat'l* Bureau Standards 48, 99 (1952). (22) R. S. Dudley, R. Elliot, W* H, McFadden and L. W* Sheraelt, J. Chem. Phys* 2S. 535 (1955). (2 3) A* G* Funk, J* C. Giddings, C, J* Christensen and H. Eyring, J. Phys. Chem. fl, 1179-83 (1957). (24) H. Gerisher, Z. Electrochem. 61. 276-80 (1957). (2 5) A. G* Funk, D. N. Chakravarty, H. Eyring and C. J. Christensen, Z* Phys. Chem.(Frank) 15. 64-74 (1958).

Funk, Giddings, Christensen and E y r i n g 2 ^ reported that in aqueous solution a metal acquires a potential as a result of chemical reactions near the metal-solution inter­ face. When the electrode is plastically strained, a change of the interface leads to a change in the measured poten­ tial. This is generally anodic and short-lived, of the order of seconds. The potential transient introduced in a copper electrode through mechanical strain was of the order of one second duration and the potential change was ordi­ narily about 0.1 volt. These results were interpreted in terms of the film rupture method with the decay of the 20 potential, that is, the return to the original value was due to the repair of the film rupture* The previous example was explained in terms of the film rupture. However, if the applied stress were large enough to exceed the elastic limit of the copper sample, permanent damage would occur that would not be restored by repair of the film alone. This is plastic deformation, the second factor mentioned previously. Brown, Mears and Dix2^ give a good generalized theory of stress corrosion of alloys.

(26) R* H. Brown, R. Mears and E. H. Dix, ASIW- AIME, 323-39 (1944).

More specific references concerning plastic defor- P7_qo mation have been reported, * J

(27) C* J. Walton, Trans* Electrochem. Soc. 85, 239 (1944). (2 8 ) U. R. Evans and M. T. Simnad, Proc. Roy. Soc. (London) 188A. 372-92 (1947). (29) Y. Druet and P. A. Jacquet, Metaux et Corrosion 22, 139 (1947)* (30) J. J. Harwood, Corrosion £, 249-50, 290-307 (1950).

According to Harwood,plastic deformation of metals tends to increase the internal energy of the metal* 21 This shifts the electrode potential to a more negative (anodic) value. This is to be expected from the relation­ ship

A P = - nPE.

The change of free energy accompanying cold working of metals has been found to be of the order of one calorie per gram although severe cold working can result in a value as high as 15 calories per gram. Conversion of these figures to potential differences indicates a potential of the order of millivolts. This shift of potential is difficult to measure accurately because of the non-uniform distribution of stress within the crystal lattice of the metal. One of the methods of relieving stress of this type is annealing.

Druet and Jacquet2^ found that cold-worked aluminum

is 40-50 millivolts more anodic than unworked aluminum in

3 per cent sodium chloride. The effect of elastic deformation compared with film rupture and plastic deformation is considered to be negligible; however, theoretical calculations for the elastic stress voltage have been performed by Shibasaki. 31

(31) Y. Shibasaki, Bull. Pac. Eng. Yokohama Ntl. Univ. 1, 125 (1951).

These calculations were for two dimensional elastic stress; however, the above author reported in a private communicatim 22 that three dimensional stress voltage produces results of the same order of magnitude* Shibasaki reported that for electrodeposited copper in an electrolyte of 0.5M copper sulfate and 0.5M sulfuric acid the maximum stress voltage obtained was less than one microvolt* For example, when the electrodeposition stress was 5 x 105 grams per square centimeter the stress voltage _7 was 5 x 10 volts. The equation applied was as follows:

tr

stress voltage

Ed - Young's Modulus

edr = strain

q - electrochemical equivalent 3 a Poisson's ratio

dd - density

These calculations were for copper, but considering the variables in the above equation, that is, Young's Modulus, the electrochemical equivalent, Poisson's ratio, and the density, it is apparent that this would give the approximate order of magnitude of the stress voltage for any metal where no slippage or rupture of the crystal lattice occurs. A statement that this voltage is of the order of microvolts would certainly be correct. 23 Motor (Moto) Electric Potential

There have been several Investigations conducted to determine the effect of moving an electrode on Its potential.32-38

(32) S. Procopiu, 2. physik. Chem. 154. 322 (1931). (33 ) ----- • Compt. rend. 202. 1371 (1936). (3 4) ----- . j. chem. Phys. 19, 121 (1936). (35) J. F. Chittum and H. Hunt, Trans. Electrochem. Soc. 21, 207 (1937). (36) C. J. Pink and H. B. Linford, Trans. Electro­ chem. Soc. J2, 401 (1937). (37) E. Newberry and G. A. Smith, Trans. Electrochem. Soc. 22* 261 (1938). (3 8 ) H. G. Bain, Trans. Electrochem. Soc. 78. 173 (1940).

It is generally agreed that rotating an electrode in solution has a definite effect on the potential of that electrode; however, there are differences of opinion re­ garding the cause. Of the theories proposed to account for this behavior the one that seems to be gaining the most support explains the behavior in terms of the destruction of the diffusion layer in the Immediate vicinity of the rotated electrode. Pink and Linford^ reported the following example of this motor electric potential involving a pure copper 24 electrode in 0.0994 M copper sulfate, which Is shown In Table 2,

Table 2.-Motor electric potential of the copper electrode.

R.P.M. Change in pot 400 2.9 8oo 3.9 1,200 5.4 1,300 5.5 3*200 3,800 7.8 4,300 5,000 8.1§ - 3 7,200 8.6 7,600 8.9 8,000 8.9 8,400 8.9

Prom these data It can be seen that this potential change reaches a maximum value at about 8 ,0 0 0 revolutions per minute. This is generally true. Pink and Linford3^ explained this behavior in terra of the destruction of the diffusion layer, that is, as the electrode is rotated faster and faster the film layer around the electrode is disturbed, giving rise to the po­ tential change as would be predicted by the N e m s t equation. This then might be considered a special form of a concentration cell, m addition, these authors reported that the thin film of solution next to a metallic electrode 25 cannot be broken down by stirring the solution with a stirring blade or bar* Newberry and S m i t h ^ reported that of the electrodes studied (copper, silver, zinc, cadmium, carbon, lead, chromium, iron, nickel, cobalt, hydrogen, and oxygen) the greatest value for the motor electric potential was ob­ tained with a hydrogen electrode on bright platinum. This value was 0,5 volts* These authors also stated that any non-conducting but porous film covering an electrode will tend to eliminate the motor electric potential.

The Effect of a Magnetic Field

If one electrode of a cell consisting of two identical electrodes Is placed in a magnetic field, a difference in chemical potential, and therefore a differ­ ence In electromotive force, should be observed since work Is involved in the transfer of matter from one field strength to another.39 Because the forces Involved In diamagnetism are extremely weak, only cells involving chemical reactions in which a change of paramagnetic

(39) 0. E. Brand, Fh* D. dissertation, State College of Washington, (1955).

susceptibility occurs would be expected to produce a measurable change In the electromotive force by the action of attainable magnetic fields. 26 The following equation was derived which expresses the change of electromotive force produced by a magnetic field

E ...... * 2nP

E = electromotive force

tM = transference number

Xjyj = molal susceptibility

- field strength

H2 = field strength

n - number of electrons involved in the transfer

P - Paraday

This equation is simplified considerably if the field strength H2 Is nearly zero, which is true when only one electrode is placed in the field. Brand tested this equation experimentally by measuring the electromotive force produced when one electrode of a cell consisting of a positive paramagnetic ion and two Identical calomel electrodes was placed in the field of a large permanent magnet. Manganous chloride and cuprlc chloride were used as electrolytes. Mercury and mercurous chloride are diamagnetic and therefore do not affect the electromotive force of the cell. Mn4"4’, being paramagnetic, did migrate under the influence of the magnetic field, resulting in a change of the Mrx1"4, concentration. For most practical purposes the presence of a magnetic field should have no appreciable effect on the action of a galvanic cell. This was shown by Brand with 1 .2 9 M manganous chloride in a cell such as described above. In a field strength of 10,900 Gauss, the maximum values obtained were 3 .3 microvolts.

Effect of Interfaclal Free Energy

Galvanic cells may be used in principle to obtain quantitative equilibrium information on interfacial free JiO energy. Oriani^ proposed the following experiment as a method of obtaining the surface tension of solids and their

(40) R. A. Oriani, J. Electrochem. Soc. 103. 19^- 201 (1956).

temperature variation in fused salts in the region of 900°C. In the cell

Ag(s) | Ag+ , KC1, LiCl | Ag(s) , both silver electrodes are of the same composition but of different sizes. The e.m.f. of these cells, after correction for any thermoelectric force, will depend on the diameter of 28 the smaller electrode and the lnterfaclal tension of sliver. The equation applicable Is

E = ------ZPd

E = e.m.f.

* lnterfaclal free energy

Vm = molar volume

Z = electrovalence

P = Faraday

d = diameter of the smaller electrode

If, for example, the diameters of the two electrodes were 0.0625 and 0.0005 inch and the value of the lnterfaclal tension of silver (

Photovoltaic Cells

There are a large number of papers in the liter­ ature concerning photosensitivity, the photoelectric or the photovoltaic effect. The latter involves the 29 conversion of light energy to electrical energy and thus It would be possible, by Illuminating one of two similar electrodes, to create a galvanic cell between two similar photosensitive electrodes. The spectral sensitivities of metal and metal coated electrodes have been determined for copper, gold 4l and silver.

(4l) P. E. Clark and A. B. Garrett, J. Am. Chem. Soc. £1, 1805 (1939).

Clark and Garrett studied the sensitivity from 3650 to 6908 A° of copper coated with oxide, silver coated with bromide and gold. For the copper-copper oxide surface the following values were reported in a solution of 0.001N potassium bromide.

Table 3.-Photovoltaic effect on the copper-copper oxide electrode.

A° Mv 3650 41.2 4047 41,4 4358 24.0 4916 7.5 5461 3 0 .0 5780 20.0 6234 2.0 6908 15.0 30 This Is only one exaiqple of a photovoltaic cell. It does however, Illustrate the possibilities of photo­ voltaic cells and the potentials that may be expected in a copper-copper oxide cell.

Properties of Aluminum

Table 4 lists several of the properties of pure aluminum.

itp Table 4,-Properties of aluminum.

Atomic weight 26.97 grams Atomic volume 9.99 Atomic number 13 Valence Chemical equivalent ^ •99 grams Electrochemical equivalent 0.0932 milligrams Specific gravity 2.706 Melting point 658.7° C. Bolling point 2 2 7 0 .6 C.

(42) N. F. Budgen, "Aluminum and Its Alloys, " 2nd Edition, Pitman Publishing Corporation, New York, N. Y., 19^7, P. 105.

The crystal form of aluminum is the face-centered cube which characterizes most ductile metals, such as gold, lead, nickel, copper and sliver. For most practical purposes the specific gravity is taken as 2.70. Much lower values have been reported but this can be attributed to porosity appearing during casting. 31 The common valence of aluminum is plus three but under special conditions the plus one valence has been reported ^

(43) E, Raijola and A. W. Davidson, J. Am, Chem* Soc. 28, 556-9 (1956). (44) S. Shlnohara and H. Ryosuki, J, Pac. Scl., Hokkaido Uhlv. Ser. IT 4, 140-6 (1952).

Aluminum and aluminum alloys have a large number of commercial uses; an Important one Is for structural purposes. This is rather surprising considering the position of aluminum in the electromotive series.

4^ Table 5,-Partial list of the electromotive series. J

Element E°(volts) Lithium 3.045 Potassium 2.925 Barium 2.90 Sodium 2.714 Magnesium 2,37 Alumlnum 1 .6 6 2 Zinc -0.762 Copper -0.344

(45) Handbook of Chemistry and Physics, 37th Edition, 1955-1956, Chemical Rubber Publishing Company, Cleveland, Ohio.

The above table Is only a partial list to show the relative position of aluminum. The E° values listed are 32 for the oxidation reaction, as is shown in the following example: A 1

Aluminum is therefore one of the more active elements. This is also shown in the following table giving the heat effects accompanying the combination of elements and radicals, calculated in grams per equivalent of the latter, for three common elements, sodium, aluminum and iron.

^ h/^awvmiw4 ciAn +-V*a Via rt +* a#* ^bAvmia^4 av\ a ^ some sodium, aluminum and Iron compounds

Compound Heat of Formation Na20 (anhyd.) 51.0 NaCl (anhyd.) 97.9 NaCl (aq.) 96.6 Na2S0ij (aq.) 164.3 A120? (anhyd.) 63.5 AICI3 (anhyd.) 53.9 AICI3 (aq.) 79.4 A 12 (S04)3 146.5 Pe203 (anhyd.) 32.6 FeClg (anhyd•) 32.1 42.6 108.4

(46) N. J. Corson, "Aluminum and Its Alloys," D. Van Nostrand Company, New York, N. Y., 1926, p. 154. 33 It Is apparent that a combination with aluminum leads to effects far exceeding those developed in Iron. From a purely chemical standpoint no stability whatsoever against corrosive attack could be expected from aluminum. Yet metallic aluminum Is much more resistant to corrosion than iron under practically any circumstance and to some extent is even more resistant than copper* The only explanation of this Is the formation of a protective film on the surface of the aluminum. Taking into consideration the ubiquity of air and the inactivity of nitrogen, the obvious con­ clusion is that this protective film consists of alumi­ num oxide* The formation of this film proceeds much more quickly in water or water vapor and the oxide is probably hydrated to some extent. Hass^ showed by means of electron diffraction, the electron microscope, and the Drude polarization technique that an instantaneous film of oxide of 15 to 20

(47 ) a . Hass, Optik 1, 134-43 (1946).

A° Is formed on exposure of aluminum to air. The film grows afterwards and stops growing after about one month. The layer reaches a total thickness of 45 A° on smooth surfaces and has a glass-like structure. 3* High purity aluminum when immersed in pure water at 20° C. to 80° C. forms surface films which are either "amorphous" or crystalline hydrous aluminum oxides. Below a critical temperature between 60° C. and 70° C. the film growth proceeds in three stages: (l) "amorphous," (2) boehmite (V), and (3) bayerite (fi). The final film thus consists of three layers.^®

(48) R. K. Hart, Trans. Faraday Soc. 53. 1020 (1957).

Hydrous aluminum oxide in aqueous electrolytes is known to have strong adsorptive powers. This property thus enables aluminum oxide to be used in chromatographic and other adsorption columns. These adsorption powers are somewhat selective, that is, hydrous aluminum oxide is known to adsorb negative ions selectively.^"^

(49) H. Holmes and R. D. Langstabb, J. Appl. chem. £,115-24 (1952). (50) K. S. W. Sing and M. R. Harris, J. Appl. Chem. Z, 397-401 (1957). (51) R. P. Graham and A. W. Thomas, J. Am. Chem. Soc. £2* 1214 (1947). (52) M. J. Pryor, R. J. Hogan and F. B. Patton, J. Electrochem. Soc. 105. 9-17 (1958). 35 One of the contributing factors in this selectivity is the semiconducting properties of aluminum oxide. Alumi­ num oxide Is a normal (n type) semiconductor, that Is, it conducts by electrons, and the selectivity for negative ions can be explained as followsNormal semiconductors form negative ions readily since the electrons in the localized levels are situated relatively high in energy and are raised to the conducting band, where they are comparatively free, with a small input of energy. More­ over, there are anion vacancies to act as active centers and increase the value of the anion adsorption energy.

(53) D. A. Dowden, J. Chem. Soc. r 242 (1950).

N type semiconductor adsorbents have the following properties

1. Highly elevated activation energy (oxides and sulfides are high) 2. To adsorb electronegative bodies there must be Impurity centers present (in AlgO^ the excess Al+ ^ ions are the Impurity centers)

(5^) P. Algraln and C. Dugas, Z. Electrochem. 363 (1952). 36 Electrons need not arise from the surface atoms of the semiconductor but can arise from the Impurity centers some distance within the Interior of the semiconductor. There are a large number of publications concerning the rectifying properties of the alumlnum-alumlnum oxide interface. The electrolytic rectifying property is at­ tributed to a double layer caused by the adsorption of negative Ions on the surface of the oxide film and the positive charge induced on the opposite side of the boun- jiji dary layer in the metal.

(44) S. Shinohara and H. Ryosuki, J. Pac* Sci., Hokkaido Univ. Ser IX, 4,140-146 (1952).

The formation of an electronic boundary layer due to adsorption on a semiconducting surface causing a charge transfer between the adsorbate and solid has been suggested by others.55-57

(55) N. P. Mott, Proc. Roy. Soc. (A) 171. 27 (1939). (56) P. B. Weisz, J. Chem. Phys. 20, 1483 (1952).

(57 ) ------.------. 21, 1531 (1953).

Muller^® observed that when very pure aluminum

(58) W. Muller and E. I&w, Aluminum 18. 478-86, 541-4, (1936). 37 dissolves In hydrochloric acid, there are three periods. During the first there Is an induction period during which aluminum does not dissolve. Next comes a period during which the reaction velocity increases very rapidly. During the final reaction period the rate remains almost constant. An Increase in the acid concentration results In a decrease of the time required for the first two periods. Ultra pure aluminum Is considerably more resistant to corrosion than the commercial grades produced. This Is due primarily to the small percentage of Impurities, usually less than 1 per cent, in the aluminum metal. Aluminum Is attacked much more rapidly by hydrochloric acid If a small amount of copper is present.59 jf aluminum

(59) 3. Tammann and V. Boehme, Zeltsch. Anorg. Chem. 226. 32 (1935).

Is heated to 200-300°C. the copper under these conditions precipitates as CuAl2 crystals which are cathodlc In nature to the bulk aluminum and thus the aluminum Is subject to particularly rapid attack.6°

(60) 0. W. Akimov and A. S. Oleschko, Korr. Met. 3^ 125 (1935). 38

Even without the formation of Cu A12 crystals the copper would be cathodic to aluminum and the net result would be the same. An Impurity of 0.1 per cent copper in aluminum will increase the corrosion rate 1600 times. Similar situations have been reported by Akimov and

O l e s c h k o ^ 0 involving FeAl^, NiAl^, and MnAl^. Non-conducting inclusions such as silicon can also interrupt the protective film and result in an acceleration of the attack of the bulk metal. The attack occurs along the boundaries of these inclusions. Another factor that must be considered in corrosion in aqueous solution is the type of ion present in the so­ lution.^1 Probably the most significant property of such ions is their ability to penetrate the aluminum oxide film

(61) S. C. Britton and U. R. Evans, J. Chem. Soc. p. 1781 (1930). and thereby be capable of reacting with and removing the underlying metal. Chlorides, bromides, and iodides show high penetrating power; the fluoride ion, which forms complexes, penetrates less, and the sulfate ions still less. Tons with oxidizing character such as chromate and nitrate show very little penetrating power and undoubtedly serve to keep the oxide film in repair. 39 The electrochemistry of aluminum, excluding the role of aluminum in electrochemical corrosion, can be summarized rather quickly as follows: reactions involving the aluminum metal-aluminum ion system are Irreversible and because of the ease of oxidation, aluminum metal can­ not be plated from aqueous solutions. The deposition of aluminum Is therefore performed in either fused salt baths or in non-aqueous solvents. The first method, that is, fused salts cells, is up to now the only process that can be successfully and profitably applied to the commercial production of aluminum. The Hall Process Is the most familiar and widely used example of fused salt deposition of aluminum. The disadvantages of fused salt depositions are the high operating temperature, the corrosive nature of the fumes and the comparatively poor throwing power. A number of non-aqueous solvents have been used as solvents for the electrodeposition of aluminum.62-69

(62) W. A. Plotnikov, J. Russ. Phys. Chem. Soc. 1 , 466 (1902). (63) H. E. Patton, Trans. Am. Electrochem. Soc. 9 (1904). (64) L. P. Yntema and L. P. Audrieth, J. Am. Chem. Soc. 5 2, 2693 (1930). (65) H. D. Blue and P. C. Mathers, Trans. Am. Electrochem. Soc. £a, 231 (1933)* 40

(6 6 ) ------.--- 51, 339 (193*0. (67) ------.------.------52, 519 (1936). (6 8 ) L. F • Audrieth, A. Long and R. E. Edwards, J. Am. Chem. Soc. £ 8 , 428 (1936). (69) T. P. Dlrkse and H. T. Briscoe, Met. Fin. 36* 284 (193o).

Some of these are ethyl bromide, absolute ethyl alcohol, formamide, acetamide, acetonitrile, ethylene glycol, etha- nolamine, and pyridine. None of these are satisfactory for commercial application.

History of the Aluminum-Platinum Electrode Pair

Earlier work with the aluminum-platinum system was carried out by Baker and M o r r i s o n , 7° and Holland.^- These

(70) B. B. Baker and J. D. Morrison, Anal. Chem. 21, 1307 (1955). (71) R. E. Holland, Master's thesis, The Ohio State University (1956). investigators used dilute electrolytes of either acetic or benzoic acid and both found that this system, because of the sensitivity to fluoride, could be used to determine microgram quantities of fluoride. This method was applied 41 by Baker for the determination of fluoride in fuming nitric acid. 72

(72) B. B. Baker, Anal. Chem. 3 0, 1085 (1958).

The determination of fluoride in this system de­ pends on the measurement of the current flowing between the platinum and the aluminum electrodes. This current Is of the order of microamperes with microgram amounts of fluoride in the cell. Holland experienced difficulty in reproducing the current-time curve reported by Baker and Morrison. Baker and Morrison reported a curve similar to that shown in Figure 2, whereas Holland repeatedly ob­ tained a curve resembling the one shown in Figure 3. In both experiments, that is with the two shapes of curves, it was found that the determination of fluoride could be accomplished with no difficulty, m Figure 2 the current was measured after one minute and in Figure 3 the plateau value of the current was measured, m both ex­ periments a plot of the current response versus the fluoride concentration gave a straight line. This triggering effect of the small quantity of fluoride has been attributed to the extraordinary solu­ bility of the aluminum oxide (AI2O3), or more correctly hydrous aluminum oxide, coating on the electrode in Pigs, 2 and 3.— Baker and Holland type curves. CURRENT CURRENT 4 4 solutions containing fluoride ions.,J Heiman reported

(73) S. Heiman, J. Electrochem. Soc. 95. 205-25 (19^9)• that the fluoride anion has the ability to break the aluminum-oxygen bond of the hydrous aluminum oxide and thus cause the dissolution of the oxide coating. A decrease in the thickness of the passivating oxide coating on the alumi­ num electrode would increase the susceptibility of this electrode to attack and therefore a larger current flow would result. The linkage in the hydrous aluminum oxide is the AI-9-AI "ol" linkage and the dissolution could be H represented as follows:

Al-0 -A1 + nP“ AlP^3-n) + o-Al

The net reaction then would be:

A1-0-A1 + 2nP“ + H+ 2A1F, or *5 This structure of the hydrous aluminum oxide, that is the "ol" linkage, has been suggested or supported by other authors.74-76

(7*0 L. Erdey, E. F. Paulik and J. Paulik, Acta. Chim. Acad. Sci. Hung. X> 27-44, 45-56 (1955). (75) A. W. Thomas and T. H. Whitehead, J. Fhys. Chem. 27 (1931). (76) J. D. Bernal and H. D. Megow, Proc. Roy. Soc. 151A, 384 (1935).

Since the appearance of the Baker and Morrison paper there have been other published reports involving similar

usage of the aluminum electrode, Johannesson77used a rotating aluminum anode and a stationary platinum cathode to determine amperoraetrieally microgram quantities of

(77) J. K. Johannesson, Chem. and Ind. April 20, p. 480 (1957).

fluoride. A standard solution of thorium nitrate was used as the titrating agent. By using this electrode system It was possible to determine from 1 .0 to 10 micrograms of fluoride in 10 milliliters of solution. The sample solution contained 0 .1 5 milliliters of perchloric acid for each 10 mllliliteis of solution. v 46 Kabanov and Polyak^ showed that during the titration of an aluminum salt with a sodium fluoride solution using the aluminum-nichrome electrode pair the potential shift at the end point was due to the sharp

(78) B. N. Kabanov and L. Ya Polyak, Zhur. Anal* Khim. 1£, 538-44 (1958). alteration of the stationary potential of the aluminum indicator electrode. The potential of the nichrome electrode does not change. The sharp potential displacement toward more negative values is due to the acceleration of the anodic process of aluminum dissolving under the influence of the fluoride. Kabanov and Polyak reported that "chloride ions eliminate complete aluminum passivity." This indicates that chloride ions, if present in high concentrations, would interfere with the determination of fluoride. Kolthoff and Sambucetti^ reported findings quite similar to those of Johannesson. Here also a rotating aluminum electrode was used in the amperometric titration

(79) I* M. Kolthoff and C. J. Sambucetti, Com. Ed., J. Am. Chem. Soc. 81. 1516 (1959)* of fluoride. This method was successfully applied to the 47 titration of fluoride in aqueous buffers with aluminum nitrate with the aluminum acting as a pP electrode in an acid solution. Kolthoff and Sambucetti stated that the work of Baker and Morrison could not be substantiated. Considering the volume and quality of the work done in this area this statement seems unjustified. Another confirmation of the Baker and Morrison work has been reported by Howard, Weber, and Weber.®0 In

(80) G. H. Howard, A. B. Weber and C. W. Weber, A.M.A. Archives Ind. Health 12 No. 3, 355-64 (1959). this work the principle of Baker and Morrison was applied to the determination of fluoride In air. The cell and electrolyte used were the same as that used by Baker and Morrison, that Is, an electrode pair of aluminum and platinum In dilute acetic acid. These authors reported that after the determination of a calibration curve, that Is, current versus the fluoride content, a rapid, quantitative and continuous determination of fluoride could be performed with ± 7 per cent precision. It was reported that chloride present in less than 100 micrograms per liter of air did not Interfere. The previous determinations all are dependent on the reaction of fluoride Ions with an aluminum electrode. 48 The following general formula Is applicable to the re­ action:

Al+ 3 + nF~ -- > Aip3-n#

In this reaction n may have a value of 1-6. The complexlng In this system has been studied by Brosset and Orrlng x who measured the constants et an Ionic strength of 0.5 and 25°C,

(8 l) C. Brosset and J. Orrlng, Svensk. Kem. Tid. 5%, 101 (: (1943)♦

The following values were reported:

Al+3 +-p- --- ^ AlF++ Qj - 1.35 x 106

A IF*-** + F" --- ^ A1P2+ Q2 = 1*°5 x 105

A1P2+ + P" --- > A1P3 Q3 = 7.1 x 103

A1P3 +-F" --- > AlFjj” Qj* = 5.5 x 102

AlPjJ +-P“ --- > A1P5= Q5 = 43

A1F5- + P~ --- > A1P6s Qg = 3

The Q*s are expressed in molar concentrations of the appropriate species, for example,

[Air**]

Ql " [ai+3hf-] ’

It Is clear that these complexes are stable and that the concentration of free fluoride ion Is small when 49 the total aluminum and fluoride ion concentration ©re comparable. These equilibrium values for the AlFff+ and AlFg were confirmed by nuclear magnetic resonance. 82

(8 2 ) R, £. Connick and R. E, Poulson, J, Am, Chem, Soc. 5153-5157 (1957).

Additional information on these comes from the work of Brosset and Orrlng who reported a measurable slow­ ness in the formation of A1F++ in solutions containing aluminum, ferric, and fluoride ions at about 10~3 m concen­ trations, It was concluded that under these conditions the net transfer of fluoride ions from the ferric to aluminum ions was slow. This does not imply, however, that the reaction between aluminum and fluoride is slow; the slow­ ness in the previous case is attributed to the ferric ions. If, for example, fluoride were added to a solution of ferric and aluminum ions it may be concluded that the fluoride reacts rapidly with the ferric ions, but more rapidly with the aluminum Ions, The complex formation between aluminum and fluoride in aqueous solution was studied by the photometric determi­ nation of the system ferric nitrate and potassium thiocyanate.®^ The resulting FeSCMT4^ complex is decolorised

(8 3 ) K, E. Kleiner, Zhur, Obshchel Khim. 20, 221-7 (1950). 50 by fluoride and In this manner the complexes of aluminum and fluoride were studied. The conclusion drawn from this work was that the formation of A1F*"* was slow and led to a slowness of re­ action between fluoride and aluminum. It Is probable that in this work the role of the ferric ion is being overlooked and as suggested by Brosset and Orrlng the slow step is actually the exchange of fluoride between iron and aluminum. Other indications that the reaction between aluminum and fluoride is not slow are the many references to the titrations of aluminum and fluoride systems using an aluminum indicator electrode. In no instance is there any mention of a slow response in the system.

Properties of Zirconium

One of the outstanding properties of zirconium is its resistance to corrosion. The high purity arc melted metal available today has a high resistance to several corrosive solutions. On the following page. Table 7 shows some of the examples. Oxygen, nitrogen, and hydrogen influence the properties of zirconium to a very high degree. They find their way into the metal at various stages in the processing and are the cause of the major problems in producing and 51 working the metal. Oxygen is the chief offender; it forme extensive interstitial solid solutions in zirconium

QJi Table 7.-Corrosion resistance of zirconium. ^

Test Solution Temperature Corrosion Rate* # by Weight °C. Mils per Year Concentrated 69.5$ 100 0.05 HNO3 Red Fuming Nitric Room 0.73 mgs. Acid** Concentrated HC1 100 0 .1 1 Mixed Acid-84# White Room 20.5 Fuming Nitric, plus 14# H2S04 80# h2so4 35 1.2 85 # h3po4 100 71.2 20# FeCl3 35 6 .8 3

•Thousandths of an inch of surface metal removed per year from a sample size 1 x 1 x 0.040 inch. ••Gain in weight of sample in milligrams.

(84) L. B. Golden, A.S.M. Symposium on Zirconium and Zirconium Alloys, 305-326 (1953). in amounts up to about 60 atomic per cent.®^

(85) J. D. Past, Foote Prints 1^, 22 (19^0). 52 The minimum temperature for oxygen adsorption by massive zirconium has been reported to be 700° and 850°

(8 6 ) L. P. Ehrke and C. M. Slack, J. Appl* Phys. 1J, 129 (1940).

Electron diffraction studies of the film formed on zirconium at temperatures between 300-600° C. reveal the existence of the monoclinic ZrOg.^ The recommended times

(87) J. E. Hickman and E. P. Gulbransen, Anal. Chem. 20, 158 (1948). for heating zirconium without serious surface oxidation are 24 hours at 500° C., 4 hours at 700° C., and not more than 88 a few minutes above 800° C.

(8 8 ) G. L. Miller, Butterworths Scientific Publi­ cations, No. 2, London (1954).

The oxide coating on the metal increases the passivity of the zirconium. This oxide coating produced during annealing is easily removed by an acid pickling bath composed of a 1 :1 mixture of water and nitric acid plus 5 per cent hydrofluoric acid.®®*

(8 9 ) R. B. Gordon and W. J. Hurford, A.S.M. Symposium, 131 (1953)* 53 It has been reported that spontaneous electrolysis occurs In a 0*1 M perchloric acid solution with a platinum 90 and zirconium electrode* This system is also sensitive

(90) S. Megregian, Anal. Chem. 2£, 1063 (1957). to added fluoride and the current response is proportional to the fluoride concentration from 2.5 to 20.0 milligrams per liter. STATEMENT OP THE PROBLEM

The electrochemistry of aluminum has been widely studied, principally with respect to the corrosion of the metal. 9 1 ,9 2

(91) M. Whitaker, Metal Industry 8 0 . 1-40 (1952). (92) U. R. Evans, "Metallic Corrosion Passivity and Protection," Longmans, Green and Company, New York, N. Y.. Edward Arnold and Company, London, England, 19^8* pp. 247-251.

Baker and Morrison^ found that a spontaneous current was generated between an aluminum and a platinum electrode in electrolytes of acetic or benzoic acid. Addition of fluoride, even in microgram quantities, caused

(70) B. B. Baker and J. D. Morrison, Anal. Chem. 2 2, 1307 (1955). an increase in the current flow between the two electrodes. The current increase was proportional to the amount of fluoride added. 71 Holland' made a quantitative study of the corrosion of aluminum in various strengths of hydrochloric acid and

(71) R* E. Holland, Master's thesis, The Ohio State Uhiverslty (1956).

54 55 repeated the work of Baker and Morrison involving the determination of fluoride, Holland experienced difficulty in reproducing the shape of the current-time curve reported by Baker and Morrison, During the study of the current-time curve with one aluminum and one platinum electrode it was observed that the shape of the current-time curve and the magnitude of the current was dependent on the size of the aluminum anode. This led to the prediction that two similar alumi­ num electrodes would develop a potential difference if the sizes were different. This prediction was found to be correct and since no evidence of similar behavior could be found in the literature the maior portion of this re­ search was directed towards the investigation of this size phenomenon. The purpose of the investigation was twofold: first, to detennine whether the area phenomenon was real or an indirect effect of some other variable, and second, to establish the mechanism or determine the dominant fac­ tors in the mechanism. Since zirconium is also covered with a protective oxide coating and is depolarized by fluoride,^0 exploratory experiments were made with zirconium electrodes to compare their behavior with aluminum electrodes.

(90) S. Megregian, Anal. Chem. 2£, 1063 (1957). EXPERIMENTAL APPARATUS

Recorder

The current and voltage measurements in this research were made with a Sargent Recorder, Catalogue number S 72150* This Is an automatic, self balancing potentlometrlc recorder which measures current or voltage and records these variables as a function of time. This instrument is versatile, having the following full scale ranges selected by a panel control switch: 1 .2 5, 2.5, 5.0, 12.5, 2 5.0, 1 2 5.0, 2 5 0.0, 5 0 0.0, 1250.0, and 2500,0. The ranges are available in millivolts, milli- amperes or microamperes with an accuracy of 0.1 per cent or 20 microvolts, whichever Is greater. A particularly useful feature of this recorder is the synchronization switch. This makes possible the automatic synchronization of external devices with re­ cording.

Ammeter

A second current-measuring device was used for confirming that the recorder was in no way contributing to the current flow between the two aluminum electrodes. This 56 57 second instrument was a Laboratory Standard D. C. Milli- ammeter manufactured by the Sensitive Research instrument Company, Mount Vernon, New York* The minimum range avail­ able with this instrument is thirty microamperes full scale*

Reaction Cell

Two types of cells were used in this research. One, for the preliminary work, was a 100 milliliter beaker above which the electrodes were supported by a separate holder. This cell is referred to as Cell II throughout this research* The second and more elaborate cell was a special all-glass cell prepared by the glass- blowers of this department for this research. An actual size diagram of this cell (Cell I) is shown in Figure 4. A primary factor in the design of this type cell was to have a removable pair of electrodes that could be easily replaced in the same position. In this manner a pair of electrodes could be removed, cleaned, and put back in the same position and thus avoid variables caused by change of position. Another advantage of this cell was the ease of maintaining an inert atmosphere above the cell solution. The solution in these cells was stirred with either a stirring blade or a magnetic stirring bar. A 58

Pig. 4.— Actual size diagram of Cell I and an electrode holder used with this cell. 59

METALLIC ELECTRODE 1* -Li 30 OUTSIDE GLASS TUBING

INSIDE 4-1 _ 1 * 8 - 0 UTS IDE

ELECTRODE INSIDE HOLDER

L _f 60 change in the stirring mechanism did not alter the electrode behavior.

Electrodes

The cathodes used throughout this research were either smooth platinum wire or a saturated calomel electrode. The platinum electrode was prepared by forming a spiral from 18 centimeters of 20 gauge platinum wire on a 5 millimeter glass rod. The glass rod was used as a support for the wire spiral. The calomel electrode was prepared in the following manner: a layer of pure mercury was placed in the bottom of a 250 milliliter wide mouth bottle, and upon this was added a layer of mercury and calomel paste. This mercury- calomel mixture was prepared by grinding the two substances together in a mortar. After the addition of the mercury and calomel paste, about 200 milliliters of a saturated solution of potassium chloride was added. The absolute value of this cell was not determined, however, the cell was checked ©gainst two other saturated calomel electrodes. The potential difference between these cells was 0.7^ millivolts or less which indicated that the cell was reliable. To avoid a change of potential because of a change of temperature the calomel cell was mounted in the constant temperature bath. 61 The salt bridges used were prepared from saturated solutions of either potassium chloride or potassium nitrate containing four per cent agar. The aluminum electrodes were prepared from either 14 gauge B&S commercial aluminum wire (99$ Al), 20 gauge "Baker Analyzed" reagent-grade aluminum wire or 99*99 + per cent aluminum sheet or rod. The reagent-grade wire was from lot number 8386 with the following analysis:

Table 8 .-Analysis of reagent-grade aluminum wire.

Impurity Per cent Insoluble in HC1 0 .0 3 0 Nitrogen compounds (as N) 0 .0 0 1 Silicon (Si) 0 .1 0 Copper (Cu) 0 .0 2 Iron (Pe) 0.05 Manganese (Mn) 0.0005 Titanium (Ti) 0 .0 2

The wire electrodes were used In either straight lengths or spirals prepared by winding the wire around a 5 millimeter glass rod. The glass rod was removed after the preparation was completed. Of these two types the spirals were the more widely used. In addition, two other types of electrodes were prepared by Mr. Gordon Laverack for use in this research. 62 One of these wes a flat circular electrode cut from a piece of aluminum sheet. To define the surface exposure of this type of electrode the lead, one face and the edges were covered with wax. The other type of electrode was a spherical electrode prepared by turning down an aluminum rod. To avoid difficulties in making electrical contact with these electrodes a lead was left attached to the electrodes. Another advantage of this was that the lead was of the same material as the electrode and thus any galvanic current flow between two metals was avoided. The only type of zirconium electrode used was a straight length of rod. The zirconium rod available was too brittle to use otherwise. The zirconium was obtained from the Carborundum Metals Company. The metal samples used were from heat Q123 for which the following analysis was given in parts per million:

Table 9.-Analysis of zirconium rod.

Impurity P.P.M.

Cr 75 Co < 20 Fe 500 Hf 55 Mg < 20 Mn 25 Pb < 20 Ti <20 V <20 63 The leads of all aluminum and zirconium electrodes were sealed into glass tubing with wax so that a constant surface exposure was maintained. This also furnished necessary support for the electrodes and minimized the difficulty of mounting them in a fixed and rigid position. In the early part of this research Dennison "Royal Scarlet" sealing wax was used to seal the electrodes into the glass tubing (aluminum melts too low to be sealed directly in glass) but a change to Kronig's cement was made because of the ease of application. The behavior of the electrodes was not altered by changing the wax.

Solutions

The solutions used were prepared from reagent- grade chemicals with demineralized double distilled water. The pH of the acid solutions was measured with a model "G" Beckman pH meter calibrated with a saturated potassium acid tartrate solution (pH 2 3.56 at 25° C.). During the period these stock solutions were being used they were stored in the constant temperature bath at 25° to avoid any delay or error in reaching 25° C. when performing an experiment.

Water Bath

The temperature in the water bath was controlled at 25° C., - 0.2, with a Penwall bimetallic probe and a 64 600 watt a.c. heater. During the summer cooling was accomplished by a fan mounted above the bath. The reaction cell was not placed directly in the water bath because of the necessity of having the magnetic stirrer beneath the cell, instead, a one-llter crystalliz­ ing dish, with an overflow tube leading back to the larger bath, was mounted adjacent to and Just above the controlled bath. The water was circulated at 800 milliliters per minute. No temperature change occurred during the transfer of water.

Balance

A Fisher chalnomatlc, magnetically damped balance was used to make the weighings in this research. The weights were calibrated with a set of class M weights from the National Bureau of Standards. EXPERIMENTAL PROCEDURE

Unless otherwise stated the following represents the procedure involved in making a typical run:

1* The recorder was turned on about five minutes before it w s b used* If this were not done, standardization of the slide wire against the Weston cell in the recorder would be impossible*

It was advisable to have the recorder grounded, for this a ground wire was available on the a*c* plug of the recorder. The red light above the off-on switch, when "on", indicated that the ground was proper* If the light did not show red, the two way off-on switch was reversed.

2. If one or more aluminum electrodes were being used they were given a three-minute preliminary Immersion in 0.01 M hydrofluoric acid. Immediately after this they were rinsed thoroughly with distilled water to remove the hydrofluoric acid, placed in the cell and adjusted equi­ distant from the walls of the cell. To avoid any loss of time in getting the electrodes into the cell after the preliminary treatment it was advisable to have the electrodes fitted Into a holder before the hydrofluoric acid treatment* This facilitated the handling of the

65 66 electrodes and also insured that both electrodes, If two similar electrodes were being used, received the same preliminary treatment. This preliminary treatment of aluminum was not an essential factor In the behavior of the two aluminum electrodes, but It did tend to give steadier current and voltage values which was desirable for quantitative measure­ ments . The preliminary treatment of the zirconium electrodes was a five minute Immersion in a 1:1 mixture of water and nitric acid plus 5 per cent hydrofluoric acid, otherwise, the procedure was the same. The adjustment of the electrodes in the cell involved spacing them equidistant from the walls of the cell. Hiis was done by actually measuring the distance between the electrodes and the cell wall.

3. During the three-minute time lapse necessary for the preliminary treatment of the electrodes the electrolyte was transferred from the stock solution to the reaction cell. For this a series of Blue Line pipets was used to obtain an accurate measurement of the volume. In Cell I the volume was 35 milliliters and In Cell II 75 milliliters.

If the cell solution were to be outgassed, the timing of the preliminary treatment was adjusted so that 67 the preliminary treatment and outgassing were completed at the same time* Another adjustment that was made during the three minute preliminary treatment was the selection of the range and pen setting on the recorder*

4. Following the adjustment of the electrodes in the cell, the stirrer was started and the connections made to the recorder (the chart drive should be on the "brake" position). These connections were made with 14 gauge copper wire with alligator clamps soldered on the ends for easy manipulation. The red Jack on the recorder is positive and the black negative.

5. Turn the chart drive selector to the "forward" po­ sition. With the synchronization Bwitch in the circuit, the circuit to the cell was completed when the recording was started. In this manner a complete curve was ob­ tained •

6. When the effect of the presence of fluoride was studied, the measurement of the current or voltage was continued for thirty minutes to establish a nearly steady state and then the fluoride (as N&F) added. The stock fluoride solution was the same strength of electrolyte as that in the cell so the pH did not change on addition of the fluoride. The fluoride was added by means of a Blue 68 Line pipet, the size of which depended on the amount of fluoride desired (most solutions contained 10 micrograms of fluoride per milliliter). The increase in volume of the cell solution did not increase the exposed surface of the electrodes since the electrodes were sealed into glass tubing.

7. To end the run the chart drive was again returned to the ''brake" position. This broke the circuit and stopped the recorder at the same time.

8. As a precautionary measure the standardization of the recorder was usually rechecked after each experiment. RESULTS

Aluminum-PIatlnum Electrode Pair

The early part of this work was directed toward the study of the alumlnum-platlnum electrode pair and the sensitivity of this system to fluoride. Tn the initial experiments with aluminum and plati­ num In 0.2 M acetic acid, curves similar to those reported by Holland were obtained (see Pig. 3). The same type of curve was obtained with the commercial and high purity aluminum anodes thus indicating that the purity of the aluminum anode was not an influential factor in the reaction of the fluoride with the metallic aluminum. To test this point further an electrode used by Baker end Morrison (ob­ tained from B. B. Baker) was exposed, under similar experimental conditions, to microgram quantities of fluoride. The current-tlme curves obtained did not resemble the nearly exponential curves reported by Baker and Morrison but did resemble the plateau type curves reported by Holland. When the reaction cell was immersed in a constant temperature bath set at 50.0° C. the current-tlme curves tended to approach the Baker type curve.

69 70 Figure 5 illustrates current-tlme curves obtained in this research: (l) with the commercial aluminum wire anodes (9956 Al) at room temperature, (2) with the Baker anode at room temperature, and ( 3) the commercial aluminum wire anode at 50.0° C. A plot of log 1 versus t (Figure 6) for the current decay of (3) above did not give a line straight enough for the application of the Blope-intercept method of inte- grating the current-tlme curve 93J for the quantitative

(93) W. M. MacNevin and B. B. Baker, Anal. Chem. 24, 986 (1952). determination of fluoride. The success of this method Is based on the linearity of the log i versus t curve. The quantitative determination of fluoride with an aluminum anode is based on the direct relationship of the current increase with the amount of fluoride added. The Holland type curve lends itself readily to measurements such as this because of the nearly constant current values of the plateau. A series of rune was made to study the current response to varying amounts of fluoride. This calibration curve is shown in Figure 7. The experimental conditions in all these runs were reproduced as nearly as possible. To achieve these conditions the same electrodes were used and the aluminum 71

Pig. 5.— Comparison of the response to 20 micrograms of sodium fluoride of the Baker electrode at 27° C., a spiral electrode (99# Al) at 25° C., and a spiral electrode (99# Al) at 50° C. (35 mis of 0.2 M acetic acid). CURRENT (fio) ISO 100 140 120 160 40 60 SO 20 24 30 99% Al SPIRAL WIRE-FORM ELECTRODE 99% 2»C AE LCRD 27*CBAKER ELECTRODE 36 PRL ELECTRODE SPIRAL AI WIRE-FORM % 9 9 24 54 48 42 IE (MIN.)TIME

60 72 6 6 73

Pig. 6 .— Log i versus t for the current decay resulting from the addition of 20 micrograms of sodium fluoride to 35 »1b of 0.2 M acetic acid at 50° C. with the aluminum-platinum electrode pair. 74 3.6 2.8 2.0 TIME TIME (MIN.) 12 0.4 o \ 1.20 1.40 1.00 1.60 0B0 0 4 0 0.60 020 (DTl) I 601 75

Fig. 7.-"Fluoride calibration curve with the aluminum (99*99% Al)-platinum electrode pair. (35 mis 0.2 M acetic acid at 25.5° C.) 31 © O T O T © © O T" o T CM o T ro O CURRENT (n a INCREASE ON ADDITION OF F) T" © o * o 01 o rs> O O “ O No F 77 cleaned for three minutes with 0 .0 1 M hydrofluoric acid each time. Also, a new sample of 0.2 M acetic acid was used. The standard procedure involving an aluminum anode was to immerse the aluminum in 0.01 M hydrofluoric acid for three minutes prior to use in the cell. This procedure produced more reproducible current-time curves and es­ tablished a nearly steady value more rapidly then electrodes not pretreated. Figure 8 shows the residual current curves for the alum:Jnum-platinum electrode pair with and without the preliminary hydrofluoric acid treat­ ment. Addition of a soluble aluminum salt to the electro­ lyte containing the aluminum-platinum electrodes and fluoride had a pronounced effect on the current value as shown in Figure 9. This behavior was utilized by others in amperometric titrations with an aluminum indicating electrode. (See references 77, 78, 79 and 80.) in order to determine whether the fluoride reaction was limited solely to the aluminum electrode two 50 milli­ liter beakers, containing 35.0 milliliters of 0.2 M acetic acid each, were connected by a saturated potassium nitrate salt bridge (nitrate does not interfere). An aluminum anode was placed in one beaker and a platinum cathode in the other. Addition of 100 micrograms of sodium fluoride 78

Pig. 8 .— Effect of the preliminary treatment, three minutes with 0.01 M HP, on the aluminum anode In combi­ nation with a platinum cathode. No NaP added. (35 mis 0.2 M acetic acid at 27° C.) CURRENT Uia) 120 40 100 0 2 0 8 60 0 4 PRELIMINARY TREATMENT MNTS IH 0M HF .01M WITH MINUTES 3 6 IE (MIN) TIME 12 16 0 2 79 80

Pig. 9 .— Effect of a soluble aluminum salt on the galvanic current flow between the aluminum-platinura electrode pair in the presence of 50 micrograms of sodium fluoride. CURRENT (no) 100 40 80 60 20 3024 IE (MIN.) TIME 36 42 'ALUMINUM 'ALUMINUM SOLUBLE 48 t l a s 81 82 to the catholyte produced no change in the current flow between the two electrodes. Prom this it was concluded that the change in the galvanic current flow between an aluminum-platinum electrode pair on addition of fluoride was due entirely to the reaction taking place at the aluminum electrode. Inspection of the current-time curves indicated that even in the absence of fluoride a small residual current flowed between the aluminum and platinum electrodes. This suggested that the aluminum was being oxidized (dis­ solved) by Just the acetic acid. This was confirmed by immersing a weighed sample of aluminum wire in 35*0 milliliters of 0.2 M acetic acid. No electrical contact was made to the Immersed aluminum so there was no galvanic current flow to an outside circuit. A spiral weighing 2.0750 grams lost 2 .7 0 milligrams in 66 hours of immersion, a loss of 0*13 per cent. Since aluminum dissolved under conditions of no electrical contact it would certainly dissolve when in contact with a substance that was cathodic to aluminum. This accounted for the residual current in the absence of fluoride.

Aluminum-Aluminum Electrode Pair

In the course of the study of the aluminum-platinum system it was observed that the shape of the current-time 83 curves and magnitude of the current was dependent on the size of the aluminum anode In the solution. Figures 10 and 11 show the type curves obtained with a large and small anode. This indicated that a difference in the current flow at the large and small electrode existed and suggested that a trial should be made using two aluminum electrodes of the same composition and hiBtory but of unequal size to see If a potential difference existed between them. In Run No. *4, shown in Figure 12, the result of using two aluminum electrodes is shown. In this experiment the size ratio was 10:1 and a steady residual current of O.35 microamperes resulted. Addition of 20 micrograms of sodium fluoride Increased the current to 9*05 microamperes, an Increase of 8 .7 0 microamperes due to the fluoride, m this instance the small electrode was anodic (negative) and the large electrode was cathodlc (positive). A thorough search of the literature of electro­ chemistry and corrosion did not reveal any previous reports of such behavior. Since this work was completed, Greene and Fontana^ have reported the observation of a potential

(94) N. D. Greene and M. G. Fontana, Corrosion 15. *1-60 (1959). difference and current flow between a large and small area 84

Pigs. 10 and 11.— Sample current-time curves with a large and small aluminum electrode in contact with a platinum cathode in 0.2 M acetic acid. 00 CURRENT H 2 m X c © m CURRENT H s m

FIGURE II 86

Fig. 12.— Current-time curve for a 10:1 ratio of an aluminum-aluminum electrode pair in the presence of 20 micrograms of sodium fluoride, (35 mis of 0,2 M acetic acid) CURRENT (ilo) 10 4 8 2 6 832 28 TIME (MIN.) 3624 40 46 87 88 of stainless steel. However, the small area had been developed as an "artificial pit" and may therefore have differed in its surface properties from the large area. Because of the possible errors in reported electrode behavior and electrode potentials due to the unawareness of the investigators of this added variable, the major portion of this research was directed toward the study of this area phenomenon. The first objective was to study the characteristics of the potential differ­ ence between the two aluminum electrodes and to prove that this was real and not caused by some factor such as dissimilar treatment of the two electrodes. One of the first experiments performed to test the reality of this observation made use of two aluminum electrodes of equal size in the same cell. These two electrodes were prepared from the commercial wire, mounted in a holder and given the preliminary hydrofluoric acid treatment in the manner described under "Procedure." There was no current flow or potential difference between these two electrodes. Prom this it was concluded that there would be no dissimilarities in the quality of the surface developed in two aluminum electrodeB of unequal size and that the potential difference between them was not caused by different treatment of one electrode. 89 The current measurements were made with the Sargent Recorder, To see If the recorder were actually- acting as a current source and inducing the potential difference, a run was made using only a milliammeter (Standard Research Instruments Company). The current values were recorded manually. Figure 13 shows the result of this run. The same behavior was observed, although the response to 20 micrograms of sodium fluoride was less than in the experiments made with the recorder. This decrease in response was due primarily to the relatively high re­ sistance of the milliammeter. On the thirty microampere range the resistance was of the order of 10,000 ohms. From this it was concluded that the recorder was not a causative factor In the current flow between the two aluminum elec­ trodes • A run was made with two spiral aluminum electrodes of unequal size in which the size of the larger electrode was made progressively smaller by reducing the number of spirals on the larger electrode until the sizes of the two electrodes was reversed. The effect of this was zero current when the electrodes were of equal size and a rever­ sal of polarity with the size reversal. The aluminum-aluminum electrode pair was sensitive to successive additions of fluoride in the same manner as the aluminum-platinum system. The result of the two ad­ ditions of fluoride is shown In Figure 14. 90

Fig• 13. — Current-tlme curve with the current values obtained with the Standard Research instru­ ment Company milliammeter. 20 micrograms of sodium fluoride added. CURRENT

Fig. I k Effect of successive additions of fluoride on the current flow between a 2:1 aluminum- aluminum electrode pair. CURRENT (|lo) 24 28 32 TIME (MIN.) TIME 36 40 4844 93 94 Since this behavior appeared to be a function of the electrode area, a series of runs was made to study the size relationship and the response of the various size ratios to added fluoride. The fluoride added was 20 micro­ grams of sodium fluoride. Table 10 summarizes the results of the size study. The size ratio, for example 2:1, means that the two aluminum electrodes were composed of two and one spirals of wire. This is equivalent to an approximate geometric surface of 2.8 and 1.8 square centimeters re­ spectively.

Table 10.-Effect of the size relationship on the response of the aluminum-aluminum electrode pair to 20 micrograms of sodium fluoride.

Run No. Size jul a (maximum) 59 12:6 1.9 61 12:6 1.4 62 8:4 1.7 63 4:2 66 4:2 8.5 67 4:2 5.9 71 4:2 8.0 65 xl: 1* 2.8 80 4:1 7.2 82 4:1 5.5 84 2:1 11.3 85 2:1 10.0 87 2:1 10.0 88 2: 1/2 1.8 86 1: 1/2 2.2

* xl s extra large 95 These runs, with the exception of numbers 65 > 87 and 88, were made with the same pair of aluminum electrodes* The size reduction was accomplished by reducing the number of spirals on one or both electrodes. The electrodes were pretreated with hydrofluoric acid and a fresh solution of acetic acid used each time. The results of the size study indicated that the 2:1 ratio gave the maximum response to fluoride. The variation in successive runs indicated that the reproduci­ bility was poor and that other variables were involved that so far were unknown. In studying the effect of pH on the potential difference between two aluminum electrodes a series of runs was made with a 4:2 aluminum ratio. These results are summarized in Table 11 showing the response to the addition of 20 micrograms of sodium fluoride.

Table 11.-Effect of pH on the response of the aluminum-aluminum electrode pair to 20 micrograms of sodium fluoride.

EH a (maximum) 1.0 <1 2.47 2.1 2.67 2.0 2.72 6.4 (ave.) 3.6 5.0 4.5 <1 8.35 <0.1 12.0 <1 96 The results in the acid range are shown In Figure 15. The results of the pH study indicated that the potential difference between two aluminum electrodes occurred in a relatively narrow pH range and that In solutions more acid than a pH of 1.0 or more basic than a pH of 4.5 the response was practically negligible. At the pH of 1.0 it was evident from the vigorous evolution of hydrogen from both electrodes that the alumi­ num was dissolving readily in the electrolyte even in the absence of fluoride. This appeared to mask the effect of the potential difference between the two aluminum elec­ trodes. Since aluminum is an amphoteric substance these same conditions are achieved in the strongly basic so­ lutions. This does not explain the values at, for example a pH of 4.5. Aluminum or aluminum hydroxide is not ap­ preciably soluble at this pH, in fact, considerably less than in solutions with a pH of 2.7. Von Berkampf^ re­ ported that the pH for incipient precipitation of aluminum

(95) S. Von Berkampf, Z. Anal. Chem. 90. 333-5 (1932). hydroxide Is 4.15 end 10.5. In addition to the disappearance of this phenomenon in strongly acidic and basic solutions, the occurrence was 97

Pig. 15.— Effect of pH on the response of the aluminum- alumlnum electrode pair to added fluoride. CURRENT itia ) 99 related to the slight solubility of the aluminum in the electrolyte. During the course of the experiments involving the study of the area factor it was noticed that the response to fluoride was altered by directing the flow of fluoride solution from the pipet to different parts of the cell solution. Further investigation revealed that by directing the flow of added fluoride toward one electrode, that electrode acted as the anode. This was true even in the case of the large electrode (cathode) although this "forced" polarity reversal was only temporary and the normal polarity was soon restored. An example of the re­ sult of adding the fluoride to the center of the solution, in the vicinity of the anode and in the vicinity of the cathode is shown in Figure 16. This was therefore a variable that had to be con­ sidered when making current measurements related to the presence of fluoride. To avoid variations caused by this, all fluoride in subsequent work was added at the center of the cell and as close as possible to the stirring bar for the purpose of obtaining rapid, uniform stirring. The restoration of the normal polarity (the small electrode anodic) after the forced reversal of polarity by directing the fluoride to the cathode indicated that the potential difference between the two aluminum elec­ trodes was real and not the result of chance behavior. 100

Pig, 16,— Effect of adding fluoride to the center of the cell solution, to the vicinity of the anode, and to the vicinity of the cathode. (35 mis of 0.2 M acetic acid) CURRENT (M.0) 20 -4 24 H SOLUTION THE TO EOJ'NaF IE (MIN.) TIME 30 36 20TNflF TO TO 20TNflF H ANOOE THE 42 46 20rNaF TO TO 20rNaF THE CATHODE THE 54 60

101 102 In addition to the study of the size relationship between two aluminum electrodes a single set of electrodes was used to investigate the response to varying amounts of fluoride. The 2:1 ratio of the aluminum electrodes was used because the size study Indicated that this combination gave the maximum response. The same pair of electrodes and a fresh solution of acetic acid was used each time. The results, as shown in Figure 17* are for the addition of 50 to 5 micrograms of sodium fluoride. Inspection of Figure 17 shows a possible trend in the current response with a maximum near 30 micrograms of sodium fluoride. Following the addition of 5 micrograms of sodium fluoride three additional runs (9 9 , 100 and 1 0 1 ) were made with 25 micrograms of sodium fluoride added. These values did not agree with the previous points and indicated that variables other than those considered thus far were in­ volved in the sensitivity of the aluminum-aluminum system to fluoride. In an attempt to obtain reproducible results with the aluminum-aluminum system several experiments were per­ formed with the same set of electrodes and the same amount of added fluoride. All known variables, such as the rate of stirring and the position of the electrodes, were main­ tained as constant as possible. Table 12 shows the results of adding 20 micrograms of sodium fluoride to successive 103

Pig, IT.— Variation of the response of a 2:1 aluminum aluminum electrode pair to changes In added fluoride. CURRENT (tio INCREASE ON ADDITION OF 10 8 O 4 2 6 o 96 9 No. 10 o 95 / 5 9 No. o 97 9 No. o o KX) No 0 y 6 9 No 20 NoF T NO O N O o 101No. 30 99 o 94 9 No. 30 0 4 o 3 9 No O 920 2 NO. 9 Not91 104 105 runs with the 2:1 ratio of aluminum electrodes. The elec­ trodes were pretreated with hydrofluoric acid and a fresh solution of acetic acid used each time.

Table 12.-Response of the same set of 2:1 aluminum electrodes to additions of 20 micrograms of sodium fluoride.

Run No. a j l a (maximum) 122 2.4 123 6 .3 124 7.3 125 9.9 126 10.5 127 9.1 128 6.0 129 6.4

These results are shown in Figure 18 with the response plotted as a function of the number of times used. This Indicated that in addition to the previous conditions considered this system was also sensitive to use. It is conceivable that the variation of response with use is directly related to a change of the thickness of the oxide layer with use since the reactivity of the aluminum electrode is an inverse function of the thickness of the oxide layer. Continual use of the electrodes with exposure to fluoride in the preliminary treatment and to acetic acid and fluoride in the cell all contributed to the dissolution of the aluminum and a decrease of the 106

p-iff 18 — Effect of use on the response of the aluminum* aluminum electrode pair to added fluoride. 107

No. 126 No. 125

No. 127

N o 124

26 No. 128

2 o N o 129 d.

N o 122

NUMBER OF TIMES USED 108 thickness of the oxide film. This suggested that this behavior was dependent in some manner on the thickness of the oxide film and therefore in order to obtain reproducible results the thickness of the oxide film must be reproduced. The existing methods of measuring the thickness of this film such as the optical or resistivity methods are not accurate enough to be applicable. Therefore, in order to obtain reproducible results It appeared that the preliminary treatment and previous use of the electrodes must be care­ fully controlled. In order to confirm that the response varied with the thickness of the oxide layer, experiments were performed with electrodes that were exposed to extreme conditions of oxidation and reduction. For these conditions a 275° C. oven and a Vycor glass tube Inserted in a combustion fur­ nace were used. Heating the samples at 275° C. In the presence of air hastened the oxidation of aluminum and Increased the thickness of the oxide layer. The Vycor tubing was fitted with nitrogen and hydrogen Inlets and an iron-constantan thermocouple, attached to a millivolt- meter, was used to monitor the temperature changes in the tubing. The aluminum samples were heated to 3OO-3500 C. in the presence of hydrogen gas. The nitrogen gas was used to flush air out of the system before heating. This decreased the air exposure of the aluminum and was also a 109 safety precaution against the formation of the explosive oxygen-hydrogen gas mixture. Increasing the thickness of the oxide film on a pair of aluminum electrodes by heating at 275° C. reduced the response to 20 micrograras of sodium fluoride from 6.4 microamperes under normal conditions to less than one microampere, a very significant decrease. A similar pair of electrodes after exposure to hydrogen at 300-350° C* also gave less than one microampere response to 20 micrograms of sodium fluoride. It Is Im­ portant to note here that the electrodes treated with hydrogen at an elevated temperature when allowed to stand in air for 30 minutes again responded to fluoride as be­ fore . The polarity of the two aluminum electrodes was reversed by special treatment of one electrode. After heating only the small electrode In the 275° oven, it acted as the cathode. From these results it appeared that the potential difference between two aluminum electrodes depended on a critical oxide thickness on the electrodes. If the film became too thick the response was greatly reduced. The most significant factor was probably the high resistance of the abnormally thick oxide layer. Also, exposure of the electrodes to a strongly reducing atmosphere decreased 110 the response. Allowing the "reduced" electrodes to stand in air, partially restored the oxide layer and increased the response to fluoride. With the significance of the oxide thickness in mind, the curve shown in Figure 18 can now be better understood. From Run No. 122 to Run No. 126 the thickness of the oxide was being reduced continually because of the exposure to fluoride in the preliminary treatment and the fluoride added to the cell. The conditions in Run No. 126 appeared to give the maximum response to fluoride. Runs No. 126 to No. 129 were performed under the same experi­ mental conditions as No. 122 to No. 126 so a further de­ crease in the thickness of the oxide layer occurred. It appeared therefore that during the course of the reduction of the thickness of the oxide layer In Runs No. 122 to No. 129 a critical thickness was reached above and below which a lower response to added fluoride was observed. Another variable that was found to have an effect on the response of two aluminum electrodes was the presence or absence of air in the cell solution. The effect of out- gassing the cell solution was an increase in the sensi­ tivity to fluoride. The removal of air did not alter the existence of the potential difference between the two aluminum electrodes. It appeared, therefore, that the Ill presence of air (oxygen) in the solution facilitated the formation of the oxide film, making the electrodes more passive and less responsive to added fluoride. In order to determine if reproducible results could be expected from the aluminum-aluminum system, ten liters of a stock solution of 0,2 M acetic acid was pre­ pared containing 20 micrograms of fluoride per 35 milli­ liters of solution. This solution was connected by means of a siphon and overflow tube to a 50 milliliter beaker containing a magnetic stirring bar and two electrodes. An Inlet tube (3 mm. i.d.) was sealed In the center of the beaker so that the solution could be introduced and dis­ tributed in a uniform manner. It was possible with such apparatus to maintain the fluoride concentration of the solution at a constant value by a continuous replenishment of the electrolyte. Two aluminum electrodes were cleaned three minutes with 0.01 M hydrofluoric acid, placed In the cell containing only 0.2 M acetic acid and the circuit completed. When the residual current reached a nearly constant value 20 micro- grams of fluoride was added to the cell solution and at the peak value of the current response the siphon was turned on and the flow adjusted at 17-19 milliliters per minute. The solution from the siphon contained 20 micrograms of fluoride per 35 milliliters and at a flow rate of 17-19 112 milliliters per minute the fluoride content was re­ plenished every two minutes. This therefore gave a constant fluoride content in the cell solution and if it were possible to get reproducible results the current flow between the two aluminum electrodes should reach a constant value. Figure 19 shows the latter, most stable part of this run. A constant value was eventually reached although these values were not reproduced in succeding runs. In order to simplify the problem of reproducibility a platinum cathode was substituted for the aluminum cathode in the flow system. This left one aluminum electrode in the solution and it was expected that this would give more nearly constant results. The results of two successive runs, No. 251 and No. 252, are shown in Figure 20. The flow of solution was maintained between 18.8 and 19.9 milliliters per minute. In Run No. 251 the variation after 45 minutes was from 65.9 to 65.2 microamperes, a 1.1 per cent change. In Run No. 252 after 45 minutes the variation was from 63.8 to 64.5 microamperes, also a 1.1 per cent change. Using the average values the differ­ ence between Run N0.251 and Run No. 252 was 2.6 per cent. This Indicated that the platinum-aluminum elec­ trode pair if used under properly controlled conditions 113

Pig. 19.— Current response of the alumlnum-alumlnum electrode pair (2:1 ratio) to a constant fluoride concentration, 20 micrograms of fluoride per 35 mis, obtained by the flow system. CURRENT (m.0 ) 3.0 h 0 6 2 8 84 78 72 66 60 _ --- 1 --- -L ---- 1 ---- I ---- 1 ---- 1 ---- 1 ____ L_ IE (MIN) TIME X 90 ± I—— 96 102//130 //l3

3 142 136 114 115

Pig. 20.— Current response of the alumlnura-platinum electrode pair to a constant fluoride concen­ tration, 20 micrograms of fluoride per 35 rals, obtained by the flow system. TIME 1X7 could be applied to the monitoring of the fluoride content of public water supplies if the chloride and iron content of the water were not too high. It appeared that nearly constant current values are obtained for both the aluminum-aluminum and platinura- aluminum electrode pairs in the flow system in which a constant fluoride concentration was maintained. This sug­ gested that reproducible values are possible, however this has yet to be shown experimentally. It has been shown in the determination of the fluoride calibration curve with the aluminum-platinum system that a straight line relation­ ship existed between the current response and fluoride concentration. To further test this a series of runs was made under the same electrode and cell conditions with Just one aluminum electrode. The results of this series are shown in Table 13.

Table 13.-Reproducing conditions on an aluminum- platinum electrode pair and the response to 20 micrograms of fluoride. Four spirals (^*8 square centimeters) of aluminum wire.

Rian No. ju.a (maximum) 268 90.8 269 91.3 270 90.2 271 89.7 118 In these four runs the total deviation was 1.8 per cent.

Baker and Morrison^ reported that It was advisable

(70) B. B. Baker and J. D. Morrison, Anal. Chem. £Z, 1307 (1955). to determine the calibration curve and run the unknown fluoride sample in the same series of runs. They reported that the calibration curve changed "for no apparent reason" over a period of time. This change has been noticed and is shown clearly in the following runs performed four days after Runs No. 268-271 under the same experimental conditions as Runs No, 268-271. The electrodes stood in 0.2 M acetic acid when not in use following Run No. 271.

Table 14.-Reproducing conditions on an aluminum- platinum electrode pair and the response to 20 micrograms of fluoride. A continu­ ation of Runs No. 268-271.

Run No. j l l q (maximum) 272 101.2 273 100.7

It appeared that on any given day a series of runs can be performed with good reproducibility, but no agreement can be expected in later experiments. 119

In an attempt to get a reproducible thickness

of the oxide layer and therefore reproducible response

to added fluoride, several runs were made with electrodes

that were treated six minutes with 0.2 M hydrofluoric acid

and stored in a dry box for a fixed period of time. It

has been reported that treating aluminum with 0.2 M hydro­

fluoric acid for six minutes^ completely removes the

(73) S. Heiman, J. Electrochem, Soc. 95. 205-225 (1949).

oxide coating on aluminum. In this manner the oxide layer was removed from each electrode and It was hoped that ex­ posure to the controlled atmosphere In the dry box would result in the formation of the same thickness of oxide on each electrode. The results obtained were not reproducible and this approach was discontinued. Platinum is capable of being at least partially covered with an oxide film and therefore the potential of a platinum electrode is greatly dependent on the a3r (oxygen) content of the solution.96-98

(96) I. J. Orestov, Nature 136. 442 (1935). (97) M. Straumanis, A. Lugge and E. Ence, Korr. u Met. 12, 148-54 (1936). (9 8 ) B. V. Ershler, Disc. Faraday Soc. 1, 269“77 (1947). 120 Ershler also reported that platinum electrodes are capable of adsorbing anions such as acetate in so- lutlon. Because of the possible variation of the potential of the platinum electrode, the behavior of the platinum electrode was measured against the reversible, reproducible saturated calomel electrode. The calomel electrode was mounted in the constant temperature bath to maintain a constant potential. The platinum electrode is cathodic to both the calomel and aluminum electrodes and should therefore exhibit comparable behavior when connected to either of these electrodes. To avoid introducing chloride into the reaction cell a double salt bridge was used, that 1s a potassium chloride salt bridge to a saturated solution of potassium chloride and potassium nitrate and a potassium nitrate salt bridge from the mixed salt solution to the cell. Figure 21 shows the effect of just rinsing the platinum electrode in water and also immersing the plati­ num electrode in hot concentrated nitric acid (recommended by Baker) on the current flow between a smooth platinum electrode and a saturated calomel electrode. Current values were measured in these experiments, however, since all conditions were maintained constant a steady current value indicated that the potential of the platinum elec­ trode was constant. 121

Fig. 21.--Current flow between a platinum and saturated calomel electrode in 0.2 M acetic acid. The platinum was either rinsed with water or immersed in hot concentrated nitric acid before use. CURRENT Old) 0 0 1 0 6 w l U Cl ~ rw~ o 8 P RNE IH WATER WITH RINSED PI No. 284 No . 283 Pt RINSED WITH WATER WATER WITH RINSED PtNo. 283 4 2 u IN - cu HOr *ft P» IMMERSED 268 No. 36

6 6 2 t IMMERSED pt IE (MIN.) TIME OC HNO, CONC. 0 6 8 4 I in ±

o« t o h 72

o - -o o n h 3 84

96 122 123 Inspection of Figure 21 indicates that even In successive runs involving similar treatment of the plati­ num electrode reproducible results are not obtained. One condition that is known to affect the potential of e platinum electrode that was not controlled during Runs No. 283-286 was the presence of air (oxygen) dissolved in the solution. To remove the air from the cell the solution was outgassed 30 minutes with nitrogen and an atmosphere of nitrogen was maintained above the solution during the runs. In addition, the platinum electrode was stored in a 0.2 M acetic acid solution prior to use. This was a pre­ caution against a possible change of the electrode poten­ tial because of the adsorption of acetate ions. Outgassing the cell solution and storing the platinum electrode in the electrolyte gave current values much more constant and reproducible than those obtained previously. This reproducibility Is shown in Figure 22. Platinized platinum electrodes were tried but these too were subject to a change of potential with time. The Increase in current in Run No. 316 after 111 minutes was caused by removing the nitrogen atmosphere from above the cell solution. This Introduced air (oxygen) Into the cell and led to the depolarization of the platinum electrode. 124

Pig. 22.--Reproducible current flow between a platinum and saturated calomel electrode in 0.2 M acetic acid. The platinum was stored in the electrolyte before use and the solution out- gassed . 0 12 24 36 40 60 120 TIME (MIN.) 125 126 The results shown in Figure 22 indicated that reproducible values could be obtained with the platinum electrode under properly controlled conditions* There is, however, a chance of error if the conditions are not properly regulated. For this r e a s o n several runs were made using the saturated calomel electrode as the reference electrode for the aluminum anode. The calomel cell was immersed in the constant temperature bath to avoid a change of potential. The same double salt bridge used with the platinum- calomel electrodes was used with the aluminum-calomel electrode pair. With the aluminum electrode it was neces­ sary to avoid chloride contamination since chloride also depolarized the aluminum electrode. Table 15 lists the response to fluoride in several experiments performed with the aluminum (4.8 square centi­ meters) and calomel electrodes. In these runs the fluoride was added from a 10 milliliter Blue Line buret in order to get a more exact fluoride measurement.

The agreement and reproducibility of these runs was considerably better than values obtained in previous experiments. This indicated that the saturated calomel electrode was better than platinum as a reference electrode under the conditions prevailing in these experiments. 127 The aluminum-calomel electrode pair, in addition to giving a more reproducible fluoride response, was also

Table 15.-Response of the aluminum-calomel electrode pair to 20 micrograms of fluoride. All solutions were outgassed 30 minutes with nitrogen and a nitrogen atmosphere main­ tained above the cell solution.

Run No. Date jtd^s (maximum) 9 8 / 5 8 7 8 . 8 330 / 11 331 78.6 11 77.8 332 fl 333 79.0

33^ 9/ 9/58fl 77.6 335 8 3 . 6 11 78.8 336 fl 337 76.6 338 fl 79.4 349 9/24/58 79.6 350 9/25/58 8 0 . 6

useful because the effect of air (oxygen) on the aluminum electrode could be shown. Table 16 summarizes these results.

The lower response in the presence of air (oxygen) was attributed to the increased passivity of the aluminum electrode because of the ease of formation of the oxide film In the presence of dissolved air. The magnitude of the potential difference between two aluminum electrodes was determined with an aluminum pair in a 2:1 ratio. This corresponds to approximate 128 surface areas of 2.8 and 1.8 square centimeters respectively. These electrodes gave a potential difference of 18 milli­ volts. This value increased to 35 millivolts when 20

Table 16.-Effect of outgassing the cell solution on the response of the aluminum electrode to 20 micrograms of fluoride. Four spirals of aluminum.

Run No. (maximum 1 Solutions not outgassed.

295 69.6 296 69.9 297 67.8 298 66.9 299 67.8 Solutions outgassed and held under nitrogen.

330 78.4 331 78.6 332 79.0 333 79.0 334 77.3

micrograms of fluoride was added to the cell. In the absence of fluoride the current was 2.5 microamperes and in the presence of 20 micrograms of fluoride a maximum of 14.3 microamperes was observed. Fink and Linford^ stated that a solution could not be stirred fast enough with a blade or magnetic stirring bar to break down the diffusion layer around a metallic 129 electrode in the solution and thus cause a change In the potential of the electrode (motor electric effect), New­ berry and Smlth^ reported that any non-conducting but porous film covering an electrode will tend to eliminate the motor electric potential. Results obtained in this research do not support these statements. For example, the residual current and fluoride response of an aluminum-aluminum electrode pair (2:1 ratio) was altered by changing the position of the electrodes in the reaction cell. The electrodes were usually spaced equidistant from the wall of the cellt However, if one electrode were moved closer to the wall of the cell it would be subject to more vigorous stirring. Table 17 summarizes the effect of position on the residual current and fluoride response of two aluminum electrodes.

Table 17.-The effect of position on the behavior of the 2:1 aluminum-aluminum electrode pair.

Run No. Residual Response Position* current Ufa) to 20ifF“Uta)

390 A 2.5 14.3 391 B 2.6 19.6 392 C 1.2 5.6

*A - The electrodes were spaced equidistant from the wall of the cell, *B - The small electrode was placed nearer to the wall of the cell than the larger electrode. *C - The large electrode was placed nearer to the wall of the cell than the small electrode. 130 These results clearly Indicated that differential stirring does alter the behavior of two aluminum electrodes in 0.2 M acetic acid. It should be noted, however, that the polarity was the same, indicating that stirring was not a factor in the origin of the potential difference between the two aluminum electrodes. Additional information on the effect of differential stirring was obtained by placing two similar aluminum electrodes in separate compartments of an electrolytic "H" cell. These compartments each contained 60 milliliteis of 0.2 M acetic acid. Differential stirring was achieved by stirring Just one compartment of the "H" cell. Figure 23 shows the relationship between the current flow and the rate of stirring (measured with a stroboscopic light). That the ratio of comers and edges or the ratio of the freshly cut surfaces to the total electrode area was not a causative factor in the potential difference between two aluminum electrodes was shown by coating the comers and edges with wax prior to use in the cell. This did not change the polarity of the two electrodes. For example, in Run No. 403 the residual current was 1.8 microamperes and a current increase of 11 microamperes resulted when 20 micrograms of fluoride was added to the cell solution. This response agrees closely with values obtained with uncoated electrodes. 131

Pig* 2 3 *— Effect of differential stirring on the current flow between two similar aluminum electrodes. CURRENT (|io) 2.8 2.0 2.4 4 0 om- 1.6 200 0 0 4 0 0 6 R.PM. 0 0 8 O O C 1200 1400 132 133 To test further the reality of the observed potential difference and the effect of area on the magni- nitude of this potential difference four similar electrodes were prepared from successive lengths of the commercial aluminum wire. These four electrodes were mounted in a lucite electrode holder to facilitate the handling of the electrodes. The original reason for preparing four similar electrodes was to connect one electrode against two or three similar electrodes in parallel and in this way study the response of a 2:1 or a 3:1 ratio of surface areas. This approach was not satisfactory because the results of connecting two or three electrodes in parallel to a single electrode were not additive, undoubtedly because of current flow in local cells in the parallel circuit. Another difficulty in this approach was that it was not possible to maintain all electrodes the same since a difference in the amount of use was sufficient to change the behavior of two electrodes. Table 18 summarizes the results of this study using four similar electrodes designated as 1, 2, 3, and 4. In addition to the spiral and straight electrodes two other types of aluminum electrodes were used in this research, these were the spherical and disk-shaped elec- trddes prepared by Mr. Laverack. Both sets of electrodes were prepared so that the approximate exposed surface areas 134 were 2.8 and 1.8 square centimeters, the same as a 2:1 ratio of the aluminum spiral electrodes. The leads, one

Table 18.-Response to 20 micrograms of fluoride of selected combinations of four similar aluminum electrodes.

Run No. Combination Response to 20 V P” Clta)

4 1 3 (1 ) 1 vs 2,3*4 0.34 4 1 3 (1 1 ) 1 vs 3,4 0 .2 8 4 1 3 (1 1 1 ) 1 vs 4 0.04

face and the edges of the disk electrodes were covered with wax, and when used in the reaction cell these electrodes were set parallel to each other with the uncoated faces of the electrodes on the Inside. The behavior of the disk and spherical electrodes was the same as that observed with the spiral electrodes, that Is, the small electrode was the anode. The response to fluoride was also of the same order of magnitude which indicated that If the relative surface areas were the same, the geometry of the electrode was not a significant factor in the magnitude of the potential difference between two aluminum electrodes. Table 19 summarizes results obtained with the different shapes of electrodes. The potential of a large and a small aluminum electrode was determined individually against a saturated 135 calomel electrode* A double salt bridge was used to prevent chloride contamination in the reaction cell* The

Table 19.-Effect of shape on the behavior of two aluminum electrodes with relative surface areas of 1.8 and 2,8 square centimeters.

Run No. Shape Response to 20>F"fffla) 84 Wire spiral M 11.3 85 10.0 86 ft 10.0 390 tl 14.3

471 Spherical 6.9* 472 tr 11.8*

473 Disk 4.8*

♦Corrected to 35 milliliters electrodes, 2:1 aluminum commercial wire spirals, were used separately but were treated in the same manner to avoid a difference in behavior because of selective treatment of one electrode. Figure 24 shows the result of the potential measurements with time for the large and the small electrode. It was evident from Figure 24 that the small electrode was anodic (-) with respect to the larger elec­ trode although the magnitude of the potential difference varied with time and eventually appeared to reverse. These measurements were made potentiometrically and therefore no current flow occurred between the aluminum and the calomel 136

Pig, 24,— Change of potential with time of a large and small aluminum electrode in 0,2 M acetic acid. 160

(-)

2201-

LARGE ELECTRODE

> 260 o—o SMALL ELECTRODE

2801-

300

20 30 40 90 60 70 60 100 TIME (MIN.) u> 138 electrodes. It Is therefore difficult to understand why the potential of the aluminum electrode changed with time* The values In Figure 24 clarify some of the diffi­ culties encountered in trying to get reproducible results with two aluminum electrodes. The potential of the elec­ trode Is dependent on the length of time it stands in solution. From Figure 24 It Is evident that a potential difference from zero to 26 millivolts is possible. If the larger electrode were placed in the cell first, a reversal of polarity could result. This also explains why a pair of electrodes on long standing in solution, even though there was no elec­ trical connection between them, exhibited a change of polarity. The electrolyte has a pronounced effect on the potential of an aluminum electrode in solution. Four similar aluminum electrodes were prepared from 20 gauge reagent-grade wire. These electrodes were treated in the same manner and immersed individually In 0.2 M solutions of formic, acetic, propionic, and n-butyric acids. The measured pH of these solutions was 2.40, 2.83, 2.85 , and 2.80 respectively. The potential measurements were made versus the saturated calomel electrode In the same manner described previously. The steady potential values reached in the formic, acetic, propionic, and n-butyric acids were 139 312, 264, 249* and 170 millivolts respectively versus the saturated calomel electrode. Inspection of* a used spiral wire-form aluminum electrode with a microscope* indicated that the upper portion of the wire was pitted more than the lower portion of the electrode. A Bausch and Lomb microscope equipped with direct lighting was used. Pour lenses were available with magnifications of 5X, 10X, 20X, and 50X. The magni­ fication of the eyepiece was 12.5X. See Figure 25 for a diagram of the spiral wlre-form electrode. The wire form electrodes, because of the curvature of the surface which gave a limited field of focus, were not satisfactory for a study such aB this. Two disk-shaped electrodes were polished to a mirror finish, shorted, and exposed to 0.2 M acetic acid. There was no pre-treatment with 0.01 M hydro­ fluoric acid. The normal behavior was noted, that is, the small electrode was anodic with a current flow of 0.6 to 0.4 microamperes between the two electrodes. After this treatment a greater pit density was observed around the outer portion of each electrode (the perimeter was covered with wax).

♦The author wishes to thank Dr. Franklin H. Beck for his assistance in the work with the microscope. Present address, Engineering Experiment Station, The Ohio State university. 140

Pig. 2 5 .— Diagrams of the spiral and disk-shaped electrodes and the apparatus used for the measurements with the Luggin capillary. 141

ANODIC WAX WAX PORTION ANODIC PORTION

CATHODIC PORTION CATHODIC PORTION

SPIRAL ALUMINUM ELECTRODE DISK-SHAPED ALUMINUM ELECTROOE

A

LUG6IN CAPILLARY

I' — ALUMINUM SAMPLE

* ^STIRRING BAR

APPARATUS USED IN LUGGIN CAPILLARY MEASUREMENTS 142 In order to prove conclusively that selective pitting occurred on one sample of aluminum, a Luggin capillary was used to measure the potential of each portion of the electrode. In this manner the local anodes and cathodes could be located. See Figure 25 for a diagram of the apparatus used for the measurements wibh the Luggin capillary. The aluminum electrode used was the disk- shaped electrode with the lead, lower face and edges coated with wax to maintain a uniform, known exposed sur­ face area (see Figure 25). The potential measurements made with the Luggin capillary showed that the portion of the electrode near the edge was consistently 8-10 millivolts more negative than the center portion of the sample, thus proving that the area of the electrode near the wax-metal interface was anodic with respect to the remainder of the metal. Thus, In the spiral wire form electrodes used frequently through­ out this research, the upper portion would be anodic with respect to the remainder of the metallic sample. These wire form electrodes were always prepared so that the stems of both electrodes were the same in length. There­ fore, In a set of electrodes the anode areas were the same but the cathode areas were different. 143 A significant factor in the action of these electrodes was the increase in resistance of the electrode in acetic acid because of the formation of a simple or complex aluminum acetate. As this aluminum acetate compound was formed it filled the pores of the electrodes and in­ creased the resistance of the electrodes. This increase in resistance as the result of dissolving aluminum in acetic acid was shown by measuring the resistance, with a conductivity bridge, of an acetic acid solution containing a piece of metallic aluminum. The indicator electrode used was platinum so the change of resistance should be due almost entirely to the change in solution as a result of the dissolution of aluminum. See Figure 26 for the re­ sults of this measurement. Ultrasonics* were found to have a pronounced effect on the behavior of two aluminum electrodes in 0.2 M acetic acid. Application of ultrasonics caused the current flow between two electrodes to drop almost instan­ taneously to zero. The characteristics of the ultrasonics were 30-35 kilocycles, 0.9 killovolts, and 200-204 milll- amperes. The effect of the ultrasonics is shown in Figure 27.

♦The author wishes to thank Dr. Y. Shibasaki for his assistance in the work involving ultrasonics. Present address, Yokohama National University, Yokohama, Japan. 144

Fig* 26.--Change of resistance with time of a 0.2 M acetic acid solution containing 2.0 grams of aluminum metal. RESISTANCE OHMS) 64001 ✓ h 0 0 8 4 0 - 0 6 5 - 0 0 4 2 h 0 0 2 3 / / - d 0 O 4 1600 " 1600 0 0 8 0 10 140 120 100 0 8 0 6 0 4 0 2 0 r s S i i » > » i > i » < i i i ■ / ✓ / / S / / TIME(HOURS) P' _____ i . i t., — i— i — o 1*5 i ke

Pig. 27.— Effect of ultrasonics on the current flow between two aluminum electrodes in 0.2 M acetic acid. CURRENT Uio) 10 6 8 / LRSNC AE APPUED WAVES ULTRASONIC 8 6 IE (MINJ TIME 10

12 7 * 1 148

Zirconium-Platinum Electrode Pair

In this research it was verified that the platinum- zlrconlum electrode pair could be used to determine small quantities of fluoride in 0.1 M perchloric acid. Figure

28 shows a calibration curve obtained with a 2.0 centi­ meter QD 123 zirconium rod and a platinum cathode. Extra­ polation of this line to zero current gave a value of 2.6 milligrams of fluoride per liter of solution. It is interesting to note that Megregian^0 reported that he obtained no fluoride response below 2.5 milligrams per liter.

(90) S. Megregian, Anal. Chem. 2£, 1063 (1957).

In these runs and all others involving zirconium electrodes a standard method of cleaning the zirconium was adopted. This was immersion in a 1:1 nitric acid solution containing 5 per cent hydrofluoric acid. This treatment resulted in a vigorous evolution of gas and also generated heat.

Zlrconlum-Zlrconlum Electrode Pair

Two zirconium electrodes, prepared from the QD 123 rod, were sealed into glass tubing with wax leaving 1^9

Pig. 28.— Fluoride calibration curve with the zlrconium- platinum electrode pair in 0.1 M perchloric acid. CURRENT (p.0) 240 200 28Ch 320 40 20r 0 i2 80 “ | " (mg/1) F" 150 151 0.8 and 2.0 centimeters exposed to the solution. A potential difference was observed between these electrodes in 0.1 M perchloric acid and as the aluminum electrodes, the smaller electrode was the anode and the larger elec­ trode the cathode. Figure 29 shows a sample current-time curve obtained with these electrodes and also illustrates the sensitivity of this pair to added fluoride. The reality of the potential difference between two zirconium electrodes was checked by measuring the current flow between the two electrodes with a milli- ammeter (Standard Research Instruments Company). The same behavior was observed. Zirconium differed from aluminum in that It was not sensitive to repeated addition of fluoride. One of the principal differences between these two metals is the solubility of the fluorides of the metals. Aluminum forms soluble fluoride complexes and therefore the surface of the aluminum does not change appreciably, whereas zir­ conium forms a slightly soluble fluoride compound, ZrF^, which forms a coating on the electrode and decreases the sensitivity of the electrode to additional fluoride. 152

Pig. 29.— Current-time curve for two zirconium electrodes, 0.8 and 2.0 centimeters long, in 0.1 M perchloric acid. CURRENT (|La) 0 9 3 > F" TIME 24 (MIN.) 0 3 36 2 4 153 DISCUSSION

The following factors were found to influence the galvanic behavior of two aluminum electrodes: 1. Sizes 2. Size ratio 3. Size reversal 4. Amount and manner of adding a triggering agent such as fluoride 3. Forced reversal of polarity 6. Pretreatment of the electrodes 7. The pH of the solution 8. Anion of the electrolyte 9. Differential stirring 10. Length of time the electrodes stand In solution 11. Ultrasonic waves

The following factors did not Influence the galvanic behavior of two aluminum electrodes but are significant:

1. Purity 2. Shape 3. Exposed metal corners and edges 15* 155 The cause of the potential difference and current flow between two similar metallic electrodes of different size is not immediately obvious when the present concepts of electrode behavior are considered. None of the factors known to cause a potential difference between two like electrodes is applicable to this situation. In the follow­ ing section each of these conditions is discussed. Also, the possible application of these conditions to the in­ terpretation of a potential difference between two like electrodes Is considered. Prepolarization of one electrode. — No potential differ­ ence was observed between two aluminum electrodes of the same size that received the seme preliminary treatment given two aluminum electrodes of unequal size. Therefore, it was assumed that the two aluminum electrodes were similar when placed in solution. Concentration cells. — Both electrodes were placed In the same solution. No potential was applied to the two aluminum electrodes, therefore, concentration cells need not be considered. 2 - 7 Thermogalvanlc cells. ' — A difference in the temperature of the two aluminum electrodes was not possible because both electrodes were placed in the same stirred solution of 0.2 M acetic acid. 156 Gravity and centrifugal cells.^"^6 — Potential differences because of either of these effects, even though smaller than those observed between the two aluminum electrodes, are not applicable to the two aluminum electrodes. Both electrodes were stationary and at the same level. Crystallographlc effect.1^ * ^ — Tt seems highly doubtful that there would be more of one crystal face exposed on two bulk samples of metal taken from successive lengths of stock wire. Also, that the smaller one would be con­ sistently anodic. Stress cells.-*-9-31 — The most significant factor in stress cells is the rupture of the protective film on stretching. This effect is short lived, of the order of one second or less, and therefore would not apply to a set of electrodes prepared for use in this research. The second factor, plastic deformation, can account for a potential change of 40-50 millivoltB in severe cold worked aluminum and is therefore of the order of magnitude of the potential difference between two aluminum electrodes. It is doubt­ ful, however, that this is a factor. First, to achieve this potential difference only the smaller of the two electrodes would have to be cold worked. Since these elec­ trodes were prepared from successive lengths of the same wire or rod and both bent In the same manner, an appreciable difference In cold working seems highly improbable. Second, 157 one of the better methods of relieving plastic deformation is annealing of the sample. It was found that annealing two aluminum electrodes did not affect the existence of the potential difference or the polarity of the electrodes. Elastic deformation results in a change of po­ tential which is of the order of microvolts, much too small to be significant in the case of two aluminum elec­ trodes. In addition, these electrodes are not under stress when they are in use so the forces causing elastic deformation are removed and the metal returns to its original form. Motor electric potent!al,32~38 — wes shown In this research that differential stirring did alter the behavior of an aluminum electrode. The normal rate of stirring in a cell containing two aluminum electrodes was 1000 revo­

lutions per minute. Inspection of Figure 23 indicates that In order to account for a residual current of the order of one microampere between two aluminum electrodes, a difference in stirring rates of 800 revolutions per minute is necessary. The possibility of having an 80 per cent difference In stirring between two electrodes Immersed in the same solution and spaced equidistant from the wall of the cell seems highly remote. Effect of the magnetic field.39 — The effect of the magnetic field Is negligible since the change of potential 158 of a metallic electrode in a field strength of 10,000 Gauss is of the order of 1 x 10**^ volts. In addition to this, both electrodes are in the same field and therefore any effects of the magnetic field should tend to cancel. 40 Curvature of the surface. — The potential difference developed because of a difference in the curvature of surface of two electrodes would be negligible under the conditions of this research. It has been calculated that at 900° C, In fused salt cells that a potential difference of the order of microvolts should be developed. Photovoltaic cells. ^•1 — Both electrodes were subject to the same Illumination In the reaction cell and therefore no photovoltaic effects should appear. In addition, the potential difference was not altered by working in total darkness or with the cell under illumination from an ultra violet light source. The first clue leading to a possible explanation of the potential difference between two aluminum electrodes of different sizes was found while viewing under a micro­ scope a used spiral wire-form electrode. The upper portion of the electrode was pitted more than the spiral portion and thus would be anodic with respect to the bulk of the electrode. This Indicated that on two aluminum electrodes such as used frequently in this research, there are actually two anodes and two cathodes involved. Considering the 159 method of changing the size of these electrodes, that is cutting down the spiral portion, the variable is the cathode area. This agrees with the suggestion by Francis and

/ 0 4 co-workers (private communication by Greene and Fontana-' with Francis) that the potential difference between an "artificial" pit and a bulk portion of the same material is due to the difference in the ratio of the wax-metal interface to the total area of the exposed metal. The implication is that the corrosion is accelerated by crevice corrosion at the wax-raetal interface and thus that area is anodic with respect to the remainder of the sample. The difference in cathode area or a reduction of the cathode area can be considered equivalent to three different circumstances that arise during the electro­ chemical action of a cell. First, a reduction of the cathode size is equivalent to the reduction of the elec­ trode area of any ordinary cell. Second, the reduction of the cathode area can be considered equivalent to the enlargement of the anode. In either of the above circum­ stances the ratio of the cathode/anode area decreases. If the effects of the cathode reduction are applied to the enlargement of a single pit representing the anode, the absolute electrode sizes in the two systems are not 160 identical. The absolute size of a corroding system can have a pronounced effect on the polarization character­ istics of the cells.99*100 Third, the reduction of the

(99) J. N. Agar and T. P. Hoar, Disc, Faraday Soc. No. 1, 158 (19^7). (100) J. T. Waber, Corrosion, 1^, 95t (1957). effective cathode area is caused by increasing the number of pit sites on the cathode. Such area reductions in­ crease the current density of the electrode. The effect of reducing the cathode size in a model electrochemical cell can be shown schematically by means of a polarization diagram such as the one shown in Figure 30. This is an instantaneous polarization diagram and it should be emphasized that this is not necessarily a steady state diagram. It is an actual corrosion polari­ zation diagram and not an impressed current polarization diagram. Diagrams such as this can be determined in three seconds or less with the proper equipment. The activity or corrosion current represented by such a diagram, because of the polarization characteristics due to the passage of current, will change remarkably with time. These effects of time and current on the electro­ chemical behavior of a cell are mutually dependent and 161

^ E • 30•-“Polarization diagram for a process under anodic control. CATHODE

s«n

The difference in behavior of a large and small aluminum anode in 0.2 M acetic acid led to the prediction that a potential difference and current flow would be developed between two aluminum electrodes of identical composition and history but of different sizes. This prediction was found to be correct. A thorough search of the literature indicated that thiB was an original observation. None of the factors known to cause a potential difference between two like electrodes is applicable to this behavior. Also, factors such as the chemical compo­ sition or shape of the electrodes are not significant. Several factors have been found that were signifi­ cant in the behavior of two aluminum electrodes in so­ lution.

1. Size - the magnitude of the potential difference was affected by the size and size ratio of the two electrodes. 2. Size reversal - the result of continued reduction of the size of the larger electrode was zero current when the sizes were equal and flow in the opposite 168 direction when the relative sizes were reversed. Triggering agent - fluoride added to the solution produced a pronounced increase in the potential difference and current flow between two aluminum electrodes. Forced reversal of polarity - the normal polarity of the electrodes was restored after forced reversal by directing the flow of added fluoride. Oxide coating - the potential difference was decreased to a negligible value by exposing both electrodes to similar con­ ditions of oxidation or reduction. pH - the potential difference between two aluminum electrodes occurred in a relatively narrow pH range. Anion - the behavior of aluminum electrodes was dependent on the electrolyte. Pretreatment - although the occurrence of the potential difference was not de­ pendent on the pretreatment, the magni­ tude was affected. Previous use of the electrodes must also be considered. 170 9. Stirring - differential stirring did alter the magnitude of the potential difference between two aluminum electrodes, It did not, however, change the polarity, 10. Length of time in solution - the magnitude of the potential difference between two aluminum electrodes was affected by standing In the solution before the circuit was com­ pleted. 11. Ultrasonics - the potential difference between two aluminum electrodes was not observed In the presence of ultrasonic waves. 12. Comers and edges - coating the exposed metal comers and edges of a pair of aluminum elec­ trodes did not alter the behavior of the electrodes.

The potential difference between two aluminum electrodes of unequal size was explained in terms of differ­ ent cathode areas with the aid of a polarization diagram. This explanation is consistent with the factorB known to influence the potential difference and current flow between two aluminum electrodes. Two zirconium electrodes Immersed in 0.1 M perchloric acid exhibited the same behavior as two aluminum electrodes. 171 It should be noted that the potential difference between two aluminum electrodes Is not an equilibrium condition and thus does not conflict with the Laws of Thermodynamics. BIBLIOGRAPHY

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I, Richard Mac Wilson, was bora in Canton, Ohio, October 30, 1930. I received my elementary education in the public school of Harlem Springs, Ohio, and my secondary education in the public school of Carrollton, Ohio. In 19^8 I enrolled at Mount Union College In Alliance, Ohio, and received the degree of Bachelor of Science in 1952. Following graduation from Mount Union I served two years In the U. S. Army, thirteen months of which were spent in Korea. I enrolled at the Graduate School of The Ohio State University In January, 1955# and received the degree of Master of Science in December, 1956. While at Ohio State I held the following positions: Assistant in the Chemistry Department in 1955 and 1956, Assistant Instructor, Analytical Division, 1957-58, Instructor of General Chemistry (Marion Branch), 1958-59# and Research Assistant, National Institute of Health, 1958-59.

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