July, 1956
HYDROGEN OVERPOTENTIAL AT IRON CATHODES IN SODIUM HYDROXIDE SOLUTIOKS 871
Acknowledgment.-Acknowledgment is made to A. Tomichek for obtaining the ignition temperature data presented in this investigation. The authors wish to express their gratitude to D. Hart for reviewing this manuscript. Appendix A’ A
= constant = constant,
AE* AH* h H H‘
= energy of activation
a
=
activity
heat of activation Planck’s constant total heat produced by pre-ignition reaction = min. amount of heat required to raise the temp. of system to the point of ignition = = =
IC. k
= specific rate
Boltzmann’s constant transmission coefficient length of coordinate of decomposition mass of activated complex m* = partition function = gas constant AS* = entropy of activation T = absolute t>emp. t = time AHr = heat of reaction per mole magnesium dH/dt = rate of heat produced by pre-ignition reaction dQa/dt = net rate of heat gain in system dnldt = no. of moles of activated complex traveling over the potential energy barrier per unit length of the decomposition coordinate per unit time $q/dt = rate of heat dissipation = designates activated state
K 1
= = = =
i
THE E F F E C T O F SOME CORROSION INHIBITORS AND ACTIVATORS ON THE HYDROGEN OVERPOTENTIAL AT Fe CATHODES I N NaOH SOLUTIONS BY I. A. AMMAR AND S. A. AWAD Chemistry Department, Faculty of Science, Cairo University, Cairo, Egypt Received October 10, 1966
Hydrogen overpotential, 9, a t Fe cathodes is measured in pure NaOH solutions at 25”. Measurements are also carried out in 0.2 N NaOH solutions to which the following organic substances have been added: ethylamine, n-butylamine, dimethylamine, tri-n-propylamine, benzylamine, pyridine, quinoline, picric acid, p-benzoquinone and m-dinitrobenzene. The numerical increase of 11upon addition of the organic substance is ex lained on the basis of a decrease in the available surface area for hydrogen discharge. On the other hand, the numerical &crease of 7 is robably caused by the depolarization of the hydrogen evolution reaction. An alternative explanation of the results is base: on the assumption that either an attraction or a repulsion, between the adsorbed atomic hydrogen and the organic molecules, take place.
Introduction The inhibition, by organic compounds, of the corrosion of iron in acid solutions was studied by Mann, Lauer and Hu1tin.l They observed that with mono-aliphatic amines, the efficiency of corrosion inhibition increased with the length of the aliphatic chain. Furthermore, the efficiency increased as the number of radicals attached to the nitrogen increased t o three. With aromatic amines the efficiency was dependent on the nature and the size of the inhibitor. Various theories were put forward to explain the action of organic inhibitors. Thus Chappell, Roetheli and McCarthy2 concluded that the inhibition of corrosion was caused by a resistive film on the surface of iron. Bockris and Conway3 measured the hydrogen overpotential, q, at an iron cathode, in the presence and absence of corrosion inhibitors and activators, in acid solutions. They observed a rise in q toward more negative values in the presence of inhibitors. In the presence of activators q was numerically decreased. Furthermore, they observed that the direct and the indirect methods of measuring overpotential gave the same results. From this they concluded that the inhibition and activation of corrosion were directly related t o the hydrogen activation overpotential (1) C. A. Mann, B. E. Lauer and C. T. Hultin, I n d . Eng. Chem., 28, 159 (1936).
(2) E. L. Chappell, B. E. Roetheli and B. Y. McCarthy, I n d . Erie. Chem., 20, 582 (1928). (3) J. O’M. Bockris and B. Conway, THIBJOURNAL, 63,527 (1949).
rather than to the formation of a surface film. Hackerman and Makrides4 explained the effect of inhibitors on the basis of general adsorption, ie., weak physical adsorption, electrostatic attraction between the inhibitor ion and the metal and chemisorption resulting in the formation of a dative bond between the inhibitor and the metal. The aim of the present investigation is to study the effect of corrosion inhibitors and activators on the hydrogen overpotential at iron cathodes in NaOH solutions. Experimental Hydrogen oirerpotential was measured in an electrolytic cell similar to that described by Bockris.6 The cell consisted of four compartments: the anode compartment, the cathode compartment, the hydrogen electrode compartment and the inhibitor (or activator) compartment. A sintered glass disc was inserted between the first two compartments to minimize the diffusion of gaseous anodic products toward the cathode. Diffusion of atmospheric oxygen into the cell was hindered by the use of water-sealed taps and ground glass joints for the construction of the cell. Electrical contact between the cathode and the reference hydrogen electrode was made through a Luggin capillary (internal diameter 1 mm.). Hydrogen, previously purified by passing it over hot copper (450’) and then over soda lime, was introduced into the cell and was divided between the four compartments. The iron electrode, in the form of a pure iron strips sealed to glass, was introduced into the cathode compart(4) N. Hackerman and A. C. Makrides, I n d . Eng. Chem., 46, 523 (1954). (5) J. O’M. Bockris, Chem. Revs., 43, 525 (1948). (6) A ”Hilger” pure iron strip, prepared by Hilger and Co., London,
England.
I. A. AMMAR AND S. A. AWAD
872
ment through a ground glass tube sealed to the top of this compartment. Two electrodes were used: one as a electrolysis electrode and the other as a test electrode. test electrode was adjusted to touch the tip of the Luggin ca illary. !'he cell was cleaned with a mixture of pure chromic and sulfuric acids. This was followed by washin with distilled water and then with conductance water. +he electrodes, previously washed with conductance water, were then fixed in their positions in the cell. The cell was filled completely with conductance water and this water was completely displaced by pure hydrogen before the NaOH solution and the inhibitor (or the activator) were introduced. NaOH solutions were pre-electrolyzed6 for 12 hours at a.cm.-a, on the pre-electrolysis iron electrode. After pre-electrolysis, the test iron electrode was lowered into the solution and the Tafel line was traced between 10-2 and 10-6 a.cm.-2. The electrode was then drawn out of the solution and the polarizing potential was adjusted such that a current of 2 ma. (2.5 X 10-8 a.cm.-*) would flow through the electrode-solution interface when the electrode was again introduced into the solution. The electrode potential was then traced as a function of time till a constant value of q was attained. The previously de-oxy enated corrosion inhibitor or activator was then added to t8e solution in the cathode compartment t o prepare a solution of known strength. The cathode potential was again traced as a function of time till a constant value was again attained. All measurements were carried out at 25" using the direct method. The potential was measured with a valve pH meter-millivoltmeter and the current with a multirange micro milli-ammeter The current densities were calcylated using the apparent surface areas. The following amines were used: ethylamine, dimethylamine, n-butylamine, tri-n-propylamine, benzylamine, pyridine and quinoline. Pyridine and quinoline were used as liquids. The other amines were used as hydrochlorides. The hydrochlorides were twice crystallized before the solution was prepared. The effect of the addition of picric acid, p-benzoquinone and m-dinitrobenzene on q also was studied. The f i s t two were dissolved in water and the third was dissolved in ethyl alcohol.
Fii
.
045.
045
Results Figure 1 shows three typical Tafel lines for hydrogen overpotential a t iron cathodes in pure 0.2, 0.5 and 1.0 N aq. NaOH solutions a t 25". Each of these three Tafel lines is the mean of six individual lines. It is clear from this figure that a t the low current density range, dissolution of iron causes the Tafel lines to become parallel to the log c.d. axis a t potentials negative with respect to the potential of the reversible hydrogen electrode. The Tafel line slope varies between 0.120 and 0.130
li
0.50 .
I
111
n
0.0
:
I
6
5
4
3
2
Log c.d. Fig. 1.-I, 0.2 N NaOH, b = 0.125 v.; 11, 0.5 N NaOH, b = 0.120 v.; 111, 1.0 N NaOH, b = 0.130 v.
Figure 2 shows the effect of the addition of ethylamine and n-butylamine on the steady-state value of q at 2.5 X in 0.2 N NaOH solution. Thus, the numerical increase of in the case of
.
TTol. 60
July, 1956
HYDROGEN OVERPOTENTIAL
.4T IRON CATHODES IN SODIUM
HYDROXIDE SOLUTIONS
873
in the case of 0.1 44 nbutylamine than in the case of 0.1 M ethylamine (cf. Fig. 2) may be ascribed to a stronger adsorption in the former than in the latter case. Similar observations have been reported by Swearingen and Schram.lO The increase in the efficiency of corro0 4 X sion inhibition in the case of n-propylamine than in the case of ethylamine has been attributed by the 03- 04. above authors to a stronger adsorption. in the case of 03 n-propylamine than in the case of ethylamine. The $ 02. 0 3 results obtained for benzylamine, pyridine and 02r quinoline (Fig. 4) may also be explained on the 01. 0 2 basis of a decrease in the available surface area for 01. 30 O l hydrogen discharge. 00It has been pointed out by Hackerman and Sudbury” that amines are incapable of undergoing 111 II 1 10 20 30 40 50 60 70 SO 90 cathodic reduction under conditions existing in Time (minutes). corrosion. For this reason it is difficult to explain Fig. 5.-I, 0.01 M picric acid; 11, 0.01 M p-benzoquinone; the decrease of q in the case of dimethylamine and 11, 0.01 ill m-dinitrobenzene. tri-n-propylamine on the assumption that such The effect of the concentration, C, of corrosion amines depolarize the hydrogen evolution reaction. inhibitors or activators on q is given in Table I. On the other hand the results of Fig. 5 may be atThus, for amines, 9 numerically increases with in- tributed to the depolarization of the hydrogen crease of concentration. For corrosion activators evolution reaction by reducible substances such as such as picric acid 9 numerically decreases with picric acid, p-benzoquinone and m-dinitrobenzene. The fact that in some cases (cf. Fig. 3, 4 and 5) increase of concentration. the decrease of q after the addition of the organic TABLEI substance is followed by an increase of q cannot be Substance Concn., moles/l. t) (mv.) explained on the basis given above. In the followEthylamine 10-1 417 ing discussion an alternative explanation of the 10-2 382 results obtained in the present investigation is 10-9 350 given. Dimethylamine 10-1 375 At the steady state, before the organic sub2 x 10-2 355 stance is added, the cathode surface is covered to a 10-a 3 13 certain extent with adsorbed atomic hydrogen. Upon addition of the organic substance, part of n-But ylamine 10-1 500 this substance is transferred to the electrode sur10-8 398 face by diffusion. Assuming for simplicity that the 10-3 293 adsorbed hydrogen is then surrounded by organic Picric acid 10-2 180 molecules, also adsorbed on the surface, one of the 10-3 325 following two processes may take place: (1) a Discussion7 repulsion between the adsorbed atomic hydrogen The mechanism of the cathodic hydrogen evolu- and the organic molecules, or (ii) an attraction, or tion at iron electrodes in pure NaOH solutions is even a chemical reaction, between the above two probably governed by a rate-determining slow dis- entities. If the heat evolved for the adsorption charge step. This conclusion is based upon the of one atom of hydrogen on a bare metal surface is value of the Tafel line slope (cf. Fig. 1). The facts designated by e, the heat of adsorption E‘, also for that the Tafel line slope lies between 0.120 and one atom of hydrogen, on a surface partly covered 0.130 v. and that only one slope is observed in the with the organic substance is given by e’ = e f SBE linear logarithmic section of the line indicates a (1) rate-determining slow discharge process.* I n alka- where S is the number of adsorption sites, occupied line solutions, it is probable that the discharge takes by the organic molecules, in the near neighborplace f r x n water molecules.s It has been sug- hood of the site occupied by adsorbed atomic hygested by Rockris and Conmay3 that the inhibition drogen, 6 is the surface coverage and E is the interand activation of corrosion are cansed by changes action energy per molecule of the organic substance. in the activation overpotential. The reason for the The plus and minus signs refer to attraction and refact that some substances numerically increase q pulsion interactions, respectively. For the adwhile some others decrease 9 is not clear. sorption of one gram atom of hydrogen, equation I The numerical increase of q upon addition of some becomes organic substances may be attributed to a decrease A = e’N = eN f SONE (2) in the available surface area for hydrogen discharge, where N is Avogadro’s number. thus increasing the current density. The fact that (10) L. E. Swearingen and A. Schram, TRIEJOURNAL,68, 180 (1951). (7) The European convention of sign of potential is used. (8) J. O’M. Bockris and E. C. Potter, J . EEectrochem. SOC.,99, 169 (11) N. Hackerman and J. D. Sudbury, J . Electrochen. Soc.. 97, 106 (1952). (1950). (9) J. O’hf. Bockris and E. C. Potter, J . Chem. Phys.. 20, 614 (12) S. Glasstone, K. Laidler and FI. E y i n g , “The Theory of Rate (1952). Processes,” McCraw-Hill Book Co., New York, N. Y.,1941, p. 364. the decrease in q is 100 and 200 mv., respectively. The dotted line in Fig. 5 indicates that q becomes positive between 30 and 38 minutes.
-
q is numerically greater
’
h
,
HUBERT T.HENDERSON AND GEORGE RICHARD HILL
874
Vol. 60
The relation between q and the heat of adsorption, A , has been given as13
with time. The fact that (dqldt) is not always linear (cf. Figs. 2, 3, 4 and 5 ) after the addition of the organic substance may be attributed to the (dv/dA) = ( 1 / 4 (3) change of V with time. As the rate of diffusion dewhere cy is the transfer coefficient. From (2) and creases with time, (dqldt) numerically decreases ( 5 ) one gets until q is almost constant a t a comparatively large eN SONE time of polarization. S C (4) = OF On the basis of the above discussion, the results where C is an integration constant. E and S may obtained for ethylamine, n-butylamine (Fig. 2), be taken as constants for a constant surface struc- pyridine and quinoline (Fig. 4) indicate a repulture, and E is a constant for each organic substance. sion between the adsorbed hydrogen and the organic molecules. On the other hand, the results Differentiating q with respect to 8 one gets for tri-n-propylamine (Fig. 3) and p-benzoquinone SNE (dvjde) = f (Fig. 5 ) indicate an attraction between the organic ( 5 ) ffF substance and the adsorbed hydrogen. Although The increase of e with time (after the addition of the assumption of attraction and repulsion may the organic substance) is directly proportional to explain some of the results obtained in this inthe rate, V , of the diffusion of this substance to the vestigation, yet it fails to account for the fact that electrode surface, thus in some cases (cf. Figs. 3, 4 and 5 ) the numerical decrease of q observed after the addition of the (dO/dt) = C‘V (6) where C’ is a proportionality constant. From ( 5 ) organic substance is followed by a numerical increase of q a t a comparatively large time of polarizaand (6) one gets tion. Similar observations have been reported by SNE (dvjdt) = f 7 C’V Bockris and Conway3 for nitrobenzene. (7) The time effect may be that caused soIely by Equation 7 indicates that for an attraction, or a the displacement of hydrogen atoms (or some other chemical reaction, between the adsorbed hydrogen sorbed material) by the inhibitor. With activators and the organic substance, ie., when E is positive q the displacement could occur either simultaneously increases with time. In other words q numerically or prior to the depolarization, but the final step in decreases with time. On the other hand a repul- each case is adsorption of the organic substance. sion interaction results in a numerical increase of q The authors wish t o express their thanks to Prof. A. R. Tourky for helpful discussion. (13) P.Ruetschi and P. Delahay, J. Chem. Phys., 2 3 , 195 (1955).
‘
A KINETIC STUDY OF METHYL CHLORIDE COMBUSTION BY HUBERT T. HENDERSON AND GEORGE RICHARD HILL Department of Fuel Technology, University of Utah, Salt Lake City, Utah Received October 91,1966
A study of the burning of methyl chloride has been carried out by the tube method of Gerstein, Levine and Wong.1 Burning velocity data have been taken in air and in oxygen. The burning velocity and limit data for methyl chloride are compared with methane and methyl alcohol and with some other chlorinated hydrocarbons. I n addition, the influence of tube diameter for the Gerstein, et al., tube method has been studied and found to be much the same as for the CowardHartwells method. A new generalized procedure for calculating flame front areas has been used.
Introduction The work presented here is a part of an over-all program to determine the effect of substituent groups in the fuel molecule on the rate of burning. For this work the simplest member of the paraffin series, methane, has been chosen as a reference. The effect of the substitution of an OH group for an H atom in methane has been studied and previously r e p ~ r t e d . ~ .The ~ work reported here concerns itself with the effects of substituting a C1 for an H atom in methane. It is believed that (1) M. Gerstein, 0.Levine and E. L. Wong, J. Am. Chem. Soc., 78, 418 (1951). (2) H. F. Coward and F. J. Hartwell, J. Chem. SOC.,Pt. 2, 2676 (1932). (3) W. H. Wiser and G. R. Hill, “The Kinetics of Combustion of Methyl Alcohol,” Tech. Rpt. No. 4, Air Force Combustion Contract No. A F 33 (038) (20839), University of Utah, 1952. (4) W. H. Wiser and G. R. Hill, “A Kinetic Comparison of the Combustion of Methyl Alcohol and Methane,” Fifth Symposium on Cornbustion, Williams and Wilkins, Baltimore, Md., 1955.
such studies as these will assist in gaining a hr! ter understanding of the mechanism of hydroc- rt,on burning.
Experimental Apparatus and Procedure The apparatus used for determining burr:: ~g velocitiep has been patterned after Gerstein, etaZ.1 Pyrex combmtion tubes, open a t both ends, were used in a horizontal position. Mixtures were ignited in all cases by a methyl chlorideoxygen flame seated on a porous glass plug burner tip. Burning mixtures were prepared by a partial pressure method and thoroughly mixed prior to introduction into the combustion tube. The combustion tubes were fitted with the appropriate orifices t o give stable, constant burning rates. Observed flame velocities were measured by means of balanced photomultiplier tubes placed behind collimating slits spaced 15.22 cm. apart near the end opposite the point of ignition. The time between impulses from these tubes as the flame front passed gave the flame velocity, UO. The flame fronts were photographed by means of an electronically actuated Speedgraphic camera a t a shutter speed of second. Flame areas, Ar, were calculated by a method which will be discussed.
