Therms
mics of
J
Processes P. A. JOHNSON AND A. L. BABB Cniaersity of Washington, Seattle, W'uslr.
ANY investigations have been made on the relation be-
cial steel in dilute potassium chloride solution as determined by both gravimetric methods and calculation of the corrosion rate from Equation 1. In addition, the concepts of the thermodynamics of irreversible processes were applied to the elect~rochemical reactions occurring on the surface of a met'al corroding in solution. From these considerations it appears that derivation of Equation 1 requires that there be no coupling between the several electrochemical reactions occurring on the electrode. A morc general equation for the corrosion current is proposcd for the case in which the c-oupling effect is not negligible.
tween cathodic and anodic polarization curves and electrochemical corrosion. Particular emphasis has been placed on the corrosion of iron because of its commercial importance, dating back to the work of Evans and Hoar (5),who demonstrated that the corrosion loss is directly related to the corrosion current. Streicher ( 1 0 ) and hlears and Brown ( 1 , 6 ) discussed the significance of the discontinuities in polalixation curves in relatioil t o
EXPEK13lENTAL PROCEDURE
CIRCUlT FOR ELECTRICAL hIEASURFMENTS. -kS the discontinuities in the relatioilship between the applied external current and the potential of the electrode behaving as a galvanic couple est,ablish the values of the desired cathodic and anodic currents used in Equat,ion 1, potential measurements should eliminate in so far as possible any resistive components. \Vhile techniqucs for this t,ype of measurement have been proposed by several investigators, the method used was adapted from one proposed by Holler (4).
I
I /
u
i
K
The basic circuit used is shorn in Figure 1, where the upper left arm of the Wheatstone bridge contains a cell of internal resistance 1' and e.m.f. Ed. The fixed resistances, Q and D,are equal, and the variable resistance, X , is set equal to r by varying X until the opening and closing of key K cause no change in deflect,ion of the galvanometer, I ' ~ ,which mill be deflected because of the e.m.f., Ed. K i t h the resistances balanced, .TT8 is then varied until the galvanometer shows no deflection. Consideration of the electrical relations under these conditions s h o w that E,] = 27',.
I-i-iq
~
This circuit is simply modified to measure the potential of L: single electrode behaving as a galvanic couple by substituhg the assembly shown in Figure 2 for the combination of Ed and T in the upper left. arm of t'he bridge. 3f is the metal electrode; S is a reference half cell, a saturat,ed calomel electrode in this study; and C is an inert carbon electrode. The potential of the metal electrode relative to the calomel half cell can then be readily determined from the measurements demribed above. With the usual potentiometer measurement, the potential of electrode LIT would consist of it,s actual potential, ITg, phi? an I R drop due to current flow, where R is the effective resistance between M and S. Xe explained more completely by Holler ( 4 ) , the bridge circuit eliminates the I R drop. It is posjible that' varying current flow will alter the over-all resistance between S and M , but fortunately such changes will not significantly affect the results, as the value sought is the current a t which there is a discontinuity in the current-potential relationship and the actual value of the potential a t this point is not of immediate consequence. Thus different values of the variable resistance, X,cause the indicated potential-current discontinuity t,o occur at, different measured values of the potential but the measured current values remain essentially constant. The total current, I , was measured with a microammeter accurate t o within 2%. The applied pot'ential, V,, was read to &0.0005 volt,; as only relative valuesof this potential were needed, the absolute accuracy was not, required. The galvanometer
___f
Q
D
Figure 1. Wheatstone Bridge Circuit
the corrosion loss of metals. The following algebraic relation between the corrosion current, io,and the magnitudes of the external currents, I*, I,, a t the discontinuities in the respective rathodic and anodic polarization curves was developed analytically by Pearson (7):
I, and I, are commonly designated as the cathodic and anodic protection currents, respectively. Holler ( 5 )developed the same relation by trigonometric means from a plot of the currentpotential relations of a galvanic couple under the application of external direct current, assuming the current-potential relations t o be linear. I n both of these developments of Equation 1, the potential of the anode or cathode a t the point where the anodic or cathodic reaction becomes zero under the application of external direct current was assumed to be equal to the open circuit potential of the respective electrodes. This study was made to compare the corrosion rates of commer518
I N D U S T R I A L A N D ENGINEERING CHEMISTRY
March 1954
had a period of 2.85 seconds, with a scnsitivity of 3.8 pa. per mm. To balance the bridge accurately, a morc sensitive galvanometer would be needed, but relative potential readings only were needed here, The fixed resistances, Q and D, were 100,000 ohms matched to within 0.5%.
I
0 Figure 2.
Arrangement of Cell
GRAVIMETRIC JJTEIGHT LOSS hIEASLJRERIENTS. CorrOSiOIl measurements were made on samples of mild stml plate approsimately 0.033 inch thick cut to specimen size of approximately 1.3 X 1.8 inches. To ensure a constant area availablc for corrosive attack, the upper one third of each specimen was coated with glyptal. Tho specimens were then immersed in solutions of approximately 0.05,V potassium chloride to a point above the edge of the glyptal coating. The solutions were maintained a t 25" i 0.1" C. by immersing the beakers containing the specimens in H, water bath. The beakers were exposed to normal room atmosphere and the slight solution losses by evaporation were compensated by additions of water as necessary. Prior to immersion the specimens were ground with an enicry wheel and polished with emery cloth to remove all oxide and leave a uniform surface. The specimens were then washed with carbon tetrachloride and weighed on an analytical balance. St the completion of the exposure period, the oxide waa removed by treatment in a solution of approximately 0.065 molal citric acid m-hile a cathodic protection current of 0.02 ampere was applied. This method had been successfully used previously (S), and tests on clean iron specimens showed it did not remove over 0.1 nig. of iron. CURREST-POTENTIAL MEASCRE 's. The rclationships between the applied current and the potential u-eve obtained with the previously described electrical circuit,. Therc does not a,ppear to be any singularly correct method for obtaining polarization curves. The most consistont data appeared to result from application of equal current increments a t evenly spaced time intervals, and an arbitrary interval of 1 minute was used in these tests between potential readings and application of current increments. The size of the current increment was adjusted to be corisist,ent with t.he magnitude and rapidity of polarization. The results of the polarization measurements were plotted in the form of current-potential curves as shown in Figure 3. In general, discontinuities were evident, but in some of the anodic polarization curves which were flatter, it was difficult to locate the discont,inuity definitely. An aid to locating this point is the fact that it occurs approximately a t the reversible potential of each electrode reaction. While the potential measurements were not absolute, they indicated the region in which the discontinuity was to be expected. The corrosion current, io, was then calculated for each observation by means of Equation 1. Both ioand Z, were plotted against time of corrosion and the average values of these variables during the corrosion period were determined from these plots. The weight loss was then calculated using a value of 2.89 X gram per coulomb as the, electrochemical equiva.lent of iron. The value of Z, was used to calculate weight loss, even though it was apparent it would yield a high result, in order to evaluate a ratio of i o / Z p for the over-all corrosion loss based only on the observed values of Zp and the measured weight loss without regard t'o any assumed relation between i o and I,. This was done to
519
detc.iniine the constancy of the 74/Zprelation, which IS of concern
111:ictical
DISCUSSION OF RESULTS
Data on the measured and calculated weight loss for nine steel specimens tested for periods of 2 to 3 days are summarized in Table I The ratio of zo/Zpshown in column five is determined from the value of 2. calculated from Equation 1. The ratio in column nine is obtained by using a value of z4 calculated from the measured weight loss. Table I shows that thc weight losses calculated from ? o did not differ by over 9% from the measured weight loss, the average deviation being 5.0%. Precise agreement between the calculated and measured dataonweight lossis not to be expected for the following reasons: The polarization curves are not linear, as was assumed in the derivation of Equation 1; the precise graphical location of the discontinuities in the polarization curves is difficult, particularly where the change in slope a t the discontinuity is small; the measured weight loss data reflect all the .i.ariation in the corrosion rate of the specimen, whereas the calculated data depended on the instantaneous values of the corrosion rate a t the times of observation; and although individual specimens were prepared for testing in the same manner, inherent differences in surface conditions may have influenced the corrosion rate of one specimen as compared to another,
0
a
I
0
-400
-
I
-
I
-
I
CURRENT MICROAMPERES Figure 3.
Polarization Curves
Tablc I also shows that the ratios of i o / Z p based on iovalues ohtained from the measured weight loss range between 0.66 and 0.82 with an average value of 0.73. These ratios were obtained independently of any fixed relation between io, I p , and Iq, and, from a practical viewpoint, would permit the approximation of corrosion rates from cathodic polarization data alone for a given environment. I n certain environments the constancy of this ratio may yield better practical values for the weight loss than those calculated by use of Equation 1; however, this was not observed in this investigation.
INDUSTRIAL AND ENGINEERING CHEMISTRY
520
TABLE I. DATAON CORROSION RATESA N D WEIGHTLoss STEELIK POTASSIUM CHLORIDE SOLUTIONS Corrosion Time, Hours 8 12 33 54 54 9 28 50 54 9 28 50 53.5 1 20 45 55 68 72 11 26 48 61 4 11 23 32 46 53 55 4 9 23 33 46 51 56 3 10 24 33 46 53 56 4 10 24 33 46 51 55 6
W t . Loss IPS
I!A
io,
pa.
Ira.
pa.
475 480 625 580
1500 1200 1600 1200
361 343 440 391
0.76 0.71 0.70 0.67
680 440 620
1160 429 1100 314 1300 420
0.63 0.71 0.68
710 575 470
1300 460 0 . 6 5 1500 416 0 . 7 2 1600 360 0 . 7 7
620 480 410 430 390
1300 420
0.68
1950 339 1550 337
0.83 0.78
530 390 370
1200 368 1360 303 800 253
0.78 0.68
i*/Ip“
0.69
575 2080 451 510 1680 391 470 1400 362 470 1160 335 515 1360 374 450 1360 338
0.78 0.77 0.75 0.71 0.73 0.73
490 3000 630 1580 460 1540 490 1310 490 1340 400 1400
0.71 0.77 0.73 0.73 0.78
490 600 375 375 380 290
1750 1360 1280 1400 1340 1540
450 354 357 359 311
383 0 . 7 8 416 0 . 6 9 290 0 . 7 7 296 0 . 7 9 296 0 . 7 8 244 0 . 8 4
FOR
from io, Mg.
Measured wt. Loss, Mg.
% Dev.
22.2
21.7
f.3
0 69
21.8
21.7
fO.5
0 69
22.7
21.4
i6.1
0 66
28.0
26.1
f7.3
0.75
17.2
15.9
f8.2
0.70
21.6
22.7
-4.8
0.78
io/lpb
Vol. 46, No. 3
For both of these phenomena it is noted that the flux is linear in the force causing the phenomenon. Khen two or more of these phenomena occur simultaneously, they interfere and cause new effects. An example is thermal diffusion resulting from the combined effects of ordinary diffusion and heat conduction. For this case a term proportional to the temperature gradient may be added to the right-hand side of Fick’s law, resulting in the statement that flow of matter can iesult from both a concentration gradient (ordinary diffusion) and a temperature gradient (thermal diffusion). ,411 of these relations define certain phenomenological coefficients L,such as heat conductivity and ordinary and thermal diffusion coefficients. [These irreversible phenomena are described by relations that ale “phenomenological” in the sense that they may be experimentally verified, but are not a part of the comprehensive theory of irreversible processes (S)]. These irreversible phenomena may be expressed by phenomenological relations of the following lorm: (5) Equation 2 is applicable where two forces are involved, which, in the diffusion example, would be the temperature and concentration gradients that give rise to the flow of matter. The reciprocal relation for the flow of heat would be: Jz = Ln X I
22.2
23.5
-5.5
0.80
20.5
18.8
f9.0
0.69
18.6
18.9
-1.6
0.82
+ Lz X,
(6)
The coefficients L I and ~ LZL are related to interference or crossphenomena and represent a “coupling” of heat flow with matter flow in the diffusion example. In accordance with Onsager’s fundamental theorem, if a proper choice is made for the fluxes, Jl,J z , and the forces, XI, XZ,then LIZ = Lzl. The proper choice of the forces is such that when each flux, J j , is multiplied by the appropriate force, Xi, the sum of these products is equal to the rate of creation of entropy in unit volume of the system. The rate of entropy creation is then given by:
i o values obtained from Equation 1.
b i o values
obtained from measured weight loss.
where T is the absolute temperature
The law of conservation
of energy always holds, but entropy is not conserved in irreverEXTENSION OF THEORY
Because the corrosion of a single metal electrode in solution is an irreversibk electrochemical process, it would be of interest to apply the principles of the thermodynamics of irreversible processes to the electrode reactions, to determine if possible the range of application of Equation 1. I n order to present more clearly the relationships for the electrochemical phenomena of interest, the basic concepts of the thermodynamics of irreversible processes a~ developed by Prigogine ( 8 ) , deGroot ( d ) , and others are given here. THEORIES OF IRREVERSIBLE P~OCESSES. It is known that a large number of phenomenological laws exist which describe irreversible processes in the form of proportionalities. Examples are Fourier’s law relating heat flow and temperature gradient and Fick’s law relating the flow of matter of a component in a mixture and its concentration gradient. These irreversible phenomena may be expressed by relations of the form
J = L X
(2)
where J is a “flux” or “velocity,” L is a scalar quantity of the nature of a conductance, and X is a thermodynamic “force” or “affinity.” For the specific examples of heat flow and diffusion, these relationa may be expressed as:
JH = k grad T
(3)
J;M = D grad G‘
(4)
aible processes. There is a net production of entropy, and the equation of continuity for entropy requires a term for the rate of entropy generation in the system, which is either zero or positive in accordance with the second law of thermodynamics. The validity of taking the phenomenological laws as linear relations between the fluxes and forces must be proved experimentally. In general, however, the linear relations apply only when the system under consideration is in the neighborhood of true equilibrium. REACTIOSSAT A CORRODIXG NETALELECTRODE.Consider a corroding metal electrode on which two half reactions, designated as a and c for anodic and cathodic, respectively, are occurring simultaneously. I n the neighborhood of equilibrium the reaction velocities (or currents) may be assumed to be linear functions of both the involved affinities (or overvoltages). Following the nomenclature and basic equations presented by van Rysselberghe ( I 1 ), the relations between the velocities of the half reactions and electrochemical potentials may be expressed as follows, when L1 and Lz are the coefficients associated with the cathodic and anodic reactions, respectively:
A, + L,2 A, = LIZ Bo + Lz A,
vc = LL
(8)
Do
(9)
where vc and v,, are the velocities of the respective cathodic and anodic half reactions, the L’s are phenomenological coefficients, the A‘s are electrochemical affinities, and LIZ is consideied equal to LZl.
INDUSTRIAL AND ENGINEERING CHEMISTRY
March 1954
I n general, I = Fv where F is the Faraday and v is the reaction Ere"). velocity, and A = q F where 7 is the overvoltage, (E' E' is the polarized potential of the couple under current flow and Erev is the thermodynamically reversible potential of a given half reaction, Using these relations, Equations 8 and 9 may be espressed in terms of currents and overvoltages:
-
Considering the irreversible effects in an electrochemical reaction, the rate of entropy production per unit volume may be expressed as:
where P R is the irreversible ohmic heat generation and PT is designated as the power of polarization, the product of the current flow and the difference between the polarized and reversible potentials. The power of polarization is expressed as: Ps = E'I
-
(EJa
+ EJc) = Pcr +
Par
=
7cIc
+
?ala
2 0 (13)
For an isolated electrode which carries no net current, I = 0, I,. Under these conditions, the and I , = -IoJ since I = I , corrosion current, io, is equal to the anodic current, I,, and hence the total power of polarization from Equation 13 becomes:
+
Pr = (E, - E,) I , = (E, - E,) io
(E'
=
Par
=
F2[Lz(E'
- EJZ + L l z (E' - E,) (E' - Eo)Z+ Llz (E' - E,)
(E'
potential value where the power of polarization for a given half reaction passes through a minimum, which for half reactions that are coupled is not necessarily the thermodynamically reversible potential. Following this reasoning, therefore, it should be possible to obtain a general expression for the corrosion current, io, in terms of the reversible potentials, E, and E,, and the coefficients LI, L z ,and LIZ. Consider first the power of polariaation of the cathode reaction as given by Equation 15 and obtain the first derivative with respect to E':
dE (%)
E,, Ea
=
2PL1 (E'
- E,) + F*LH(2E' - E, - E,) (18)
Equating the first derivative to zero results in the following value of E', which makes P,, a minimum:
As it is assumed that the discontinuity in the anodic polarization curve occurs a t the value of E' given by Equation 19, the value of the applied external current will then be I,. From Equations 10 and 11 the respective cathodic and anodic currents a t this point are:
(14)
The powers of polarization for the cathodic and anodic reactions may be expressed as the product of the respective currents, as given by Equations 10 and 11, and the associated overvoltages:
P,,
521
- E,)] cathode - E,)]
(15) anode (16)
I t is noted that the cathodic current as given by Equation 20 is not zero but is finite, depending on the magnitude of LIZ. I n this interpretation, a current value different from the observed I , would be necessary to reduce the cathodic current and hence the cathodic reaction velocity to zero. Since I , is the total applied current a t the above point and is equal to the sum of I , and I,, it may be shown that:
The total power of polarization is the sum of Equations 15 and 16, and may be expressed in terms of the total current, I , as:
(LlLZ
- L212)
+ LP + 2L12) + F z (Ll + Lz + 2L12) ( E , - Ea)' 1 2
p , = FZ (L1
These relations developed for irreversible electrochemical processes may now be applied to the reactions occurring on the surface of a corroding metal electrode under the application of direct cathodic and anodic currents. CORROSION CURRENT EQUATION INCLUDING COUPLING.For a single electrode with more than one electrochemical process occurring on its surface the coupling of an electrochemical half reaction by one or several others, corresponding to the coefficients, L12, is possible, as indicated by Equations 8 and 9. In considering these reactions on a corroding steel electrode under the application of direct cathodic and anodic currents, the chemical affinities-i.e., E,,,/P-are not changed but the polarized potential, E', is varied. Consequently, there should be states of the system for which the respective powers of polarization for each half reaction as given by Equations 15 and 16 should be a minimum. The minimum values should be obtained from these equations by differentiating with respect to E', the polarized potential which changes with curient flow, and holding E, and E,, the reversible potentials, constant. When the coupling effect between the electrochemical half reactions is not assumed to be negligible, a general expression for the corrosion current, io, may be developed. In order to accomplish this, the potential of the electrode, E', where it changes in behavior from a metal electrode to a hydrogen electrode must be defined in terms of reversible potentials. It would appear reasonable to assume that this point of transition would be a t the
I n an analogous manner for the anodic reaction, the following relations may be shown to exist:
- Ea)' + F2L1z (E' - E,) (E' - Ea) (23) = 2FzL2 (E' - E,) + F2L12 (2E' - E , - Ea)
Pan = FzLz (E' dE' (%)
Ec, Ea
For a minimum value of Paw:
For the case where there is no net current flowing, ( I = 0), equating 14 and 17, results in:
There are now three independent relations as given by Equations 22, 28, and 29, in which the only variables that cannot be calculated from theoretical considerations or measured experimentally are L1, Lz, and LIZ.
522
INDUSTRIAL AND ENGINEERING CHEMISTRY
Equation 29 may be considered as a general expression for the total corrosion current flowing on a single metal electrode in solution. CORROSION CCRRENTEQCATIOX SEGLECTING COTPLISG. When L12 is considered to be zero, Equations 22, 28, and 29 mav be written as:
Solving Equations 30 and 31 for L? and 151,respectively, conibining these results with Equation 32, and recalling that I , is a negative current result in the follon ing expiession for the corrosion current:
(33) Equation 33 is the same expression for i, as previously developed in Equation 1. For this case in which LIZ is assumed to bc zero, the values of E' from Equations 19 and 25, which make t,he powers of polarization for the cathode and anode reactions a minimum, occur a t the reversible potentials of the cathode and anode reactions, respectively. Thus at the discontinuity in t,lie cathodic polarization curve n.here the esternal current is I, = I,,the corresponding potential is E' = E,, and I , = 0. Similarly, at the discontinuity in 6he anodic polarization curve, I , = I,, E' = E,, and I , = 0. X a G N i ~ r u n E OF THE COUPLISGEFFECT. It is of interest to estimate the relative magnitudes of the phenomenological coefficients. Data xere not obtained in this study to allow accurate calculat'ions to be made of the reversible potentials of t'he tTvo half reactions on t,he steel elect.rode, but for illustrative purposes some data obtained by Schwerdtfeger and McDorman ( 9 ) on an iron-copper couple in a pot,assium chloride solution will be used t o illustrate the approach to obtaining values for the coefficients. The experimental conditions under which these data were obtained were such that local corrosion on the two electrodes was negligible. The reversible potential for the anode reaction a t the iron electrode will be assumed to be -0.54 volt: the reversible potent'iai of the cathode react'ion at the copper elect,rode, -0.25 volt; the value of I,, -201 microamperes: the value of I,, 1048 microamperes; and the values of the corrosion current, io, as determined by weight loss measurements and calculated from Equation 1 were 165 and 168 pa., respectively. The reversible potentials are relative to a saturated calomel electrode. By solving Equations 22, 28, and 29 by trial and error values and signs were assigned to the coefficients as f o l l o w to obtain curLB = 39.35 X rent values in amperes: LI = 7.35 X l O - j 4 ; LIB = 0.25 X 10-14. The units of these coefficient,s are (coulombs-volts-seconds) -1 expressed in the common electrical units. The above example illustrates the method of determining the magnitude and direction of the coupling effect from this type of experimental data. This small value of L12 is to be expected where the agreement between the weight loss as calculated by use of Equation 1 and the measured value is close. In instances where the disagreement, is greater, it will probably be due to coupling effects Tvhich will be reflected in larger values of the coupling coefficient, Liz. The magnitude of L12 appears to be a measure of the applicability of Equation 1and its sign will determine whether the actual weight loss is greater or less than that predicted from Equat'ion 1. The data in Table I indicate that the weight losses as calculated from Equation 1 are greater than the measured wright losses and therefore the sign of ,512 would appear to be positive. Honever,
Vol. 46, No. 3
because of the difficulties inherent in this experimental method l d that experiments of this as previously discussed. it ~ ~ o u seem type are of limited value in determining precisely the magnitude and direction of the coupling effect. From these considerations it would appear that the nonlinear polarization curves shown in Figure 3 may result to a large ext,ent from the coupling between the electrochemical half reactions, ani1 that Equation 1 for the corrosion current, io.st,rictly app1ic.s only n.hen the coupling effect is negligible. Furt,her ~ o r l i n this direction is being conducted at this laborat,ory and it is hopctl that considerations of this nature n-ill be of value in the interpretation of corrosion phenomena. SL13131.ARY
\\-eight loss results for inxi specimens corroding in pot chloride solutions indicate that the relationship for the corrosion current of Equation 1 is accuyate withiii loyo. From the concepts of the thermodynamics of irreversible processes. the relation for the corrosion cui~rentgiven by Equation 1 and previously developed by other means is shown to result for the case where there is no coupling betivern the cathodic and anodic reactions. This coupling, holvever, is likely to occur in practice, particularly d i e n both cathodic and anodic processes occur on the same electrode. 131- considering that the discontinuities in t,he polarization curves occur where the pon-ers of polarization for the respective half reactions are a minimum, a general relat,ionship was tlereloped which should relate the rorrosion current and the cathodic and anodic protection currents even when thsre is an appreciable coupling effect. Although this relationship has not yet been conipletely studied experimentally, it appears that the coupling coefficient may in general be small for iron specimens, and its magnit,ude is in a sense a measure of the applicability of Equation 1. I t is possible, however, that E'quation 1 might not hold as well for metals other than iron, owing to an appreciable coupling effect. .4s the ratios of i o / I p for the iron specimens did not deviate more than 127, from the average value of 0.73, it would bc practical to approximate t,he corrosioii rate of steel in this environment from a measurement of the cathodic prot'ection current alone. \O\IE\C
A
=
L 4 I'UKE
electrochemical affinity
= concentration = diffusion coefficient
D
E' = polarized potential E,, E , = reversible electrode potentials of anodic and cathodic, F I I,, I , I,
= = = = = = =
I,
7Jri
= = = = = =
J,lf
h
Ll, L , LPi,L:, P, S = 1' = t
,
=
2)
=
X
= =
7
reactions, respectively the Faraday external applied cuirent anodic and cathodic currents, respectively cathodic protection current anodic protection curient corrosion current flux or velocity flux or flow of heat flux or flov of mass roefficient of thermal conduction phenomenological coefficients coupling or drag coefficients power of polarization entropy per unit volume absolute temperature time reaction velocity thermodynamic foice or affinity overvoltage
Subscripts a refers to conditions at the anode.
c refers to conditions a t the cathode.
INDUSTRIAL AND ENGINEERING CHEMISTRY
March 1954
523
,-~--, .
(4) Holier, H. L., J . Electrochem. Soc., 97, 271 (1950). (5) Ibid., p. 453. ( 6 ) Alears, R. B., and Brown, R. H., Ibid., 97,7 5 (1950).
RECEIVED for review August 17, 1953. ACCEPTEDOctober 31, 1953. Presented before the Division of Industrial and Engineering Chemistry a t the 124th Meeting of the A M E R I C A F C H E M I C A L S O C I E T Y , Chicago, 111.
4ction of Polar Organic Inhibitors
A
IN ACID DISSOLUTION OF METALS NORMAN HACKERR‘IAN AND A. C. MAKRIDES Department of Chemistry, The University of Texas, Austin, Tex.
T
HE three prerequisites of an electrochemical mechanism of corrosion are a potential differenre, a conduction path, and availability of electrode reactions for transferring charges across the metal-solution interface ( 1 4 , 18). Thus, an inhibitor may function: (1) by increasing the true ohmic resistance and (2) by interfering with the anodic, the cathodic, or both the electrochemical processes. Examples of case 1 are inhibition by the formation of an oxide film or by precipitation of a nonconducting reaction product onto the metal. Inhibitions caused by an increase in the activation hydrogen overpotential, a decrease of potential differences on the metal surface, or activation polarization of anodic dissolution are examples of case ( 2 ) . THEORIES OF CATHODIC INHIBITION
,4theory for the action of organic inhibitors was suggested by Chappell, Roetheli, and McCarthy (ti), who studied the effect of quinoline ethiodide on cathodic and anodic polarization of iron and steel in iV sulfuric acid and concluded that inhibition was cathodic. The same conclusion was reached by Mann, who proposed a comprehensive theory of inhibition by organic compounds (31). Essential features of this theory ( 6 , 31-33) are that organic inhibitors are capable of forming onium ions and accordingly exist in acid solution as cations. These are cathodically adsorbed by virtue of electrostatic attraction and thus blanket cathodic areas. The resultant film increases the interfacial resistance to passage of current by preventing hydrogen ions from reaching the surface, the nature of the cathode not actually being changed. Depending on extent of adsorption, the closeness of packing in the adsorbed film, and cross-sectional area of the molecule, various degrees of inhibition are found. Thus, according to Mann, organic inhibitors operate by mechanism 1. Evidence for this theory comes from cathodic polarization studies and from changes of the inhibitive power caused by substitution on the inhibitor. Results of measurements of film resistance ( 2 , 50, 41) are conflicting and difficult t o interpret theoretically. Machu (30) found a direct relationship between film resistance and inhibition. Bockris and Conway ( d ) , however, found a negligible film resistance and concluded that Machu’s explanation of inhibition as a resistance effect was highly improbable. A number of polarization studies have been reported (2, 6, 13, 68,34, 39, 44). In general, the major effect of inhibitors a t the current densities employed was on cathodic polarization, At high current densities and very negative potentials, as used in such
studies, the effect of inhibitors on the cathodic reaction undoubtedly becomes greater. However, these conditions are far renioved from those existing during corrosion (open circuit) and any conclusions drawn from such studies a r e likely to be erroneous (10, 26). Mann’s theory enjoyed a wide following, mainly because it conformed to the inherent notion that positively charged particles should be adsorbed on negatively charged areas. There are, however, numerous difficulties barring its acceptance. One major difficulty is presented by the change of the open circuit potential in the cathodic direction generally observed on addition of inhibitors. This shift can be explained easily if one assumes that the main effect of the inhibitor is on the anodic reaction (25). To account for this potential change on the basis of Mann’s theory, or any modification of it, the assumption must be made that the organic compound is either reducible or can depolarize the cathode (13). It is difficult, however, to see how a compound a t some fixed concentration can be simultaneously a depolarizer and an inhibitor of the same process. Furthermore, many inhibitors which give such shifts-e.g., the amines (83, %)-are incapable of undergoing cathodic reduction under conditions existing in corrosion. A second major difficulty is that both anodic and cathodic effects occur in polarization studies a t small current densities. Cavallaro and Bolognesi ( 4 ) found by polarization methods that a number of inhibitors were of a mixed type and in many cases prevalently anodic. Hackerman and Sudbury ( 2 3 ) report both anodic and cathodic effects in polarization studies with n-octylamine. Kuznetsov and Iofa (29) also report that in many cases the increase in overpotential caused by inhibitors is greater on anodic polarization. A third difficulty arises from specific effects observed with inhibitors. Sulfur-containing compounds are better inhibitors than corresponding nitrogen compounds. The theory of cathodic inhibition makes no provision for such effects, since electrostatic forces are not specific. If there is any difference in the extent to which an amine or a thiol exists as a cation in acid solution, the amine ought to exist to a greater extent as the cation, because it is the more basic. Hackerman and Cook ( 1 9 ) found that irreversible adsorption of acids, alcohols, and esters occurred on the same portion of a steel surface, while amines adsorbed irreversibly on a different portion. This is another instance of specific adsorption which indicates that other forces besides electrostatic ones are operative.
