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Inhibition of electrode processes on copper by AHT in acid solutions
INHIBITION
OF ELECTRODE AHT IN ACID J. W.SCHULTZE
PROCESSES SOLUTIONS
ON
COPPER
BY
and K. WIPPERMANN
Institut fur Physikalisohe Chemie II der Universitit Diisseldorf, UniversitHtsstraBe 1, D-4000 Dusseldorf, F.R.G. (Received 21 August
1986; in revised firm
17 November 1986)
Aktract-3-Atnino-5-heptyl-1,2,4-triazole (AHT)* is a very effective inhibitor of copper corrosion in neutral as well as acid solutions. From capacity-potential curves and their time dependence it follows that a compound layer of AHT (Cu, Cl) of 1-2 nm thickness is formed which is stable up to 0.4 V. At E > 0.45 V AHT is dissolved immediately. Current-potential curves were measured for all possible electrode reactions. The inhibitor affects the anodic oxidation of copper strongly. The cathodic oxygen reduction is influenced as well but to a lesser degree. The inhibition of the cathodic hydrogen evolution is only important in deaerated solutions and at large cathodic overvoltages. From the values and the shift of transfer coefficients we must conclude that the reaction site differs for transfer reactions of electrons, protons and copper ions. The influence of the AHT concentration is exnlained by the thickening of the layer, the conductivity of which differs for ions and electrons.
INTRODUCTION
addition of 10m3 M AHT causes’s decrease of the current density to very small values in the potential range - 0.5 to + 0.4 V. The capacity-potential curves show that this inhibition efficiency is combined with a strong decrease of the electrode capacity by a factor of almost l/10. At very high potentials only the increase of the capacity and the anodic current density indicates a desorption of the inhibitor. To investigate the electrode reactions (l)-(5) we chose corresponding systems, both without (cAHT = 0) and with inhibitor (10m5 M < cAHT 6 lo-’ M), as follows:
It has been well known for a long time that benzotriazole inhibits the copper corrosion in neutral and alkaline solutions, but it is less effective in acid solutions[lG3]. The inhibition is due to the formation of thick layers of copper-benzotriazole complexes and probably also due to the formation of copper oxide layers protected by the inhibitor. A few years ago the inhibitor 3-amino-S-alkyl-1,2,4-triazole (Trademark Texamin AT 1) was shown to have the same or even better inhibition efficiency especially in acid -Alutions[4--6]. In spite of the practical success of this inhibitor no scientific investigations of the manifold reactions which can take place in this system in neutral solutions have been carried out. Table 1 shows a summary of possible reactions (l)-(S) and the corresponding electrode potentials. Preliminary investigations in neutral solutions showed a complex overlap of anodic corrosion, oxide formation and inhibitor layer formation.
System (1) 1 N HISO+, PO, = 0 atm, 0.1 M Cu” (2) 1 N H2S04, PO, = 0 atm (3) 1 N HZS04, PO, = 0 atm, 0.1 M H202 (4) 1 N HZS04, PO, = 0.2 and 1 atm
Electrode reaction (see Table 1) (lak(W (2) (3)
Pa) + (4b)
The results of the measurements will be presented below. Further results regarding the influence of pH, anions, temperature, the comparison with benzotriazole and XPS measurements will be presented in Ref. [7].
Scheme 1. To simplify the system we carried out extended investigations in acid solutions where we could exclude the oxide formation reactions [(ld)-(lf)]. Figure 1 shows examples of current-potential and capacity-potential curves in 1 N HZSO, both with and without inhibitor. The anodic current density is due to the dissolution of copper (la), while the cathodic reaction at small overvoltages is caused by oxygen reduction (4) or at large overvoltages by hydrogen evolution (2). The
EXPERIMENTAL The electronic set-up of the electrochemical measurements is described elsewhere[S]. It consisted of a fast rise potentiostat, an X-Y recorder and a signal source and lock-in amplifier for the UC measurements. Chemicals
*Commercial inhibitor of Henkel KGaA with Trademark Texamin AT 1.
1 N sulfuric acid with a pH of 0.4 was used as electrolyte. It was prepared with p.A. HZS04 (Merck)
823
824
J. W. SCHULTZE AND K. WIPPERMANN
Table
1. Possible reactions at copper electrodes in aqueous solutions Standard potential Reactions
CWWI
(la) Cu +cu=++2e(lb) Cu *cu++e(lc) cu +*cw++.(Id) 2Cu+H~O=Cu,O+2H++-te(le) CuZO + HZ0 =2Cu0+2H++2em (If) Cu,O+ 3HZO&2Cu(OH)Z +2H+ (2) H,=2H+ +2e(3) 2HzO &H202+2Hf+2e(4a) 2H20=0, +4H+ +4e(4b) H20,==0,+2H+ +2e(5) AATw + Cu G= (Cu-AAT)+ (5b) AAT + Cu + F= (Cu-AAT) (5~) AAT + Cu2+ + (Cu-AAT)‘+
0.4
‘: E Y $r T=
Cull N H~O&,;0.2
+2e-
+ 0.336 + 0.520 +0.153 +0.471 + 0.669 + 0.747 0.000 + 1.770 + 1.229 +0.682 -
at.AHT
Cell and electrodes The electrochemical measurements were carried out in a vessel of duran glass thermostated at 25°C. As working electrode we used copper cathodically deposited on a gold wire[9]. The gold wire had a diameter of 0.4 mm and an area of 0.25 cm’. A gold foil of 4 cm2 was used as counter electrode. To check the influence of the crystal plane we used a (11 l)-oriented copper single crystal, which was kindly supplied by Professor Schaarwachter (Universitat Dortmund). The single crystal plane was mechanically polished using diamond spray (15 p, 3 p) and finally electropolished in 75% H3P04 with an anodic current density of 0.4 A cm-*. For the detection of intermediates we used a rotating split ring-disc electrode which consisted of a central copper disc (99.999 o/0Cu) of 3.4 mm diameter surrounded by 2 Pt half-rings[lO]. The collection efficiency of 1 Pt half-ring was determined to be N = 0.15 in agreement with calculations from geometry. The rrd electrode was mechanically polished with 3 p diamond spray and washed with alcohol and water before each experiment. RESULTS
0.2 0
Capacity
measurements
Fig. 1. (a) Potentiodynamic current potential curves recorded at a copper (111) single crystal in 1 N H2S04 with cAHT = 10e3 M 1-l and 0 M I- 1 respectively. Aerated solution, d&/d/dr = 5 mV s- ‘. (b) Electrode capacity under the same conditions. 1 kHz, 2 mV amplitude.
Since the electrode capacity indicates structural changes at the interface, we recorded the capacity in inhibitor solutions of various concentrations. 1N H,SO, without Cu2+ was chosen so that we could extend the measurements over a large range of electrode potentials. Systems 2 and 4 (deaerated and aerated) gave equal results. Hence we always used system 4. Further, there was no difference between single crystal surfaces and the polycrystalline Cu surface. Hence, polycrystalline material was always used. Figure 2 shows the electrode capacity measured during a slow potentiodynamic sweep under stabilized conditions (after more than 5 cycles). In 10m4 M inhibitor solution one observes already a strong decrease of electrode capacity which then becomes very small in 10e3 M and lo-* M solutions. Capacity values of indicate the formation of a very about S PF cm-’ stable inhibitor layer in the entire accessible potential range. Using the Frumkin model of parallel condensors for the adsorption of organic molecules, we estimated the coverage of AHT, 13,using the equation (1): @= (c,=.-c)/c,=.-c*~,). (1)
and three times distilled water. The solutions were either aerated by oxygen (99.999%) or de-aerated by nitrogen (99.999 %)_ The inhibitor 3-amino-5-heptyl1,2,4-triazole (AHT) was used as a hydrochloride which has a better inhibition efficiency than its sulfate form[7]. The p& value of AHT was determined as 9.35. The saturation concentration of AHT hydrochloride in water is 0.1 mol l- ‘. Correspondingly the standard free enthalpy of the solution 15.6 kJ mol- ‘. The inhibitor was kindly supplied by Henkel KGaA Dusseldorf.
Taking C,= r to be the minimum value of C for a lo-’ M AHT solution, Ca= r = Cmin = 2.5 pF cme2, we calculate (3 > 0.9 for lo- 3 M solutions and 8 z 1 for lo-’ M AHT. Such an estimation, however, is misleading, since AHT is not adsorbed reversibly under the conditions of Fig. 2, which correspond to an almost steady state of the interface. Figure 3 shows that for a fresh surface (curve 1) the capacity decreases continuously during cyclization if the desorption potential of 0.45 V is not exceeded. An almost constant value of about 5 PF cm -2 is reached in the 5th cycle. Exceeding 0.45 V, the inhibitor dissolves and the capacity increases up to 60 PF cm-’ and remains high during the cathodic cycle Sb.
-0.2 50
40 v $ 30 Y a. z 20
10
0 -1.0
0
-0.5
0.5
E(SHEIIV
Inhibition of electrode processes on copper by AHT
CullN
H2SOI,pO;0.2 at,AHT
/ /
I ; cAHT=O R-\
,
\
I
\
c-’
\
w--e
a
-0.4
I
\
,
.
825
\
-0:z
0
2’
0:2
0.4
E(SHE)IV
Fig. 2. Electrode capacity of copper in 1 N H,SO, with various concentrations of inhibitor under steady state conditions. Aerated solution &/dt = 5 mV s-l_
80
Cull
NH,S0,,p,2=
0.2 at,lO-‘M
AHT
60
0
-0.4
-0.2
0
0.2
0.4
flSHE)/V
Fig. 3. lst, 2nd and 5th potentioclynamic capacity-potential curves in 1 N H2SO4 after addition of lo-’ hi AHT. Broken line gives results of potentiostatic pulse measurements from + 0.5 V (free surfarr) to 0.3,O and - 0.5 V, respectively. To test the reversibility of the adsorption system we carried out additional potentiostatic step experiments, starting from a very positive potential, 0.5 V, where the inhibitor is des0rbe.d completely. Then we stepped the potential to the measuring potential and recorded the electrode capacity time-dependence. At f0.3 V the steady state value of Fig. 2 is approached in a short time but at lower potentials we observed a very slow decrease of the electrode capacity indicating the slow formation of the inhibitor layer. Since molecular adsorption reactions are usually fast and diffusion
limitation is to be expected for a few seconds only, an adsorption equilibrium should be maintained in few seconds. Hence, the strong deviation between the potential step experiments and the curve 1 in Fig. 3 on the one hand, and the steady state potentiodynamic sweep experiments in Fig. 2 and curve 5a in Fig. 3 on the other hand, indicate that the inhibitor layer formed in the steady state is not an adsorption layer but a stable precipitate which is formed slowly and irreversibly and does not change with electrode potential. Finally, to compare these results with a simple
826
J. W. SCHULTZEAND
adsorption system of AHT we measured the electrode capacity on mercury electrodes[GJ. For this system the lowest capacity values observed were about clearly exceeding the values observed on 8 PFcm-‘, copper electrodes. Hence, we conclude again that the AHT layer on copper electrodes cannot be described by a simple adsorption system (5a) of Table 1 but rather by equations (Sb) and (5c), respectively. More evidence for this interpretation was obtained from electrolyte exchange experiments and XPS measurements after extended rinsing of the electrode with pure water[7]. From the XPS measurements we conclude that the thickness of the inhibitor layer at the free corrosion potential is within the range of 1-2 nm. Taking the condenser formula d = DDo /C
(2)
we obtain from C = 2.5 FF cm-‘, a dielectric constant of D = 5, a film thickness of 1.8 nm, which is a reasonable value for a layer of inhibitor molecules. Measurements in sulphuric acid with AHT sulfate[7] and AHT chloride yielded smaller capacities for the chloride solution indicating a chloride catalysed formation of the inhibitor layer. Hence, we must conclude that in the system Cu/10m3 M AHT chloride, layers of 1-2 molecules are formed in the steady state, containing an unknown amount of cosorbed copper and chloride ions. Therefore, the influence of adsorption is small or negligible, and it is clear that polycrystalline electrodes and Conner single crvstals behave equally in AHT solution. _-
K. WIPPERMANN Copper
dissolution
and deposition
We carried out investigations of the anodic copper dissolution and cathodic copper deposition in system 1 using a rotating ring-disk electrode in 0.1 M Cur+. The upper part of Fig. 4 demonstrates the dependence of the disk electrode current density on the electrode potential. In solutions without inhibitor we see the transfer controlled reaction with large transfer currents as seen in Fig. 4a between 0.2 V and 0.35 V. The lower part of Fig. 4 shows the corresponding ring current at a potential of0.4 V, indicating the formation of Cu+. The ring current density increases with increasing potential almost parallel to the equilibrium concentration of Cu + calculated from the Nernst equation for reaction (1 b). In solutions with 1O-3 M AHT the current density at the disk as well as at the ring is almost 0 in the range from 0.2 V to 0.4 V (Fig. 4a). At E > 0.47 V the rapid rise in the disk current indicates desorption of the inhibitor which allows anodic dissolution of copper. This current decreased more slowly for the cathodic sweep. Simultaneously Cu+ is indicated at the ring electrode. Below 0.2 V a small amount of copper deposition takes place with a current density of about without production of Cu+. The slow 1 mA cm-‘, hysteresis observed for the cathodic current can be attributed to the production of a fresh surface during the copper deposition which allows a faster transfer of copper ions in the following anodic sweep, up to 0.2 V. The existence of this hysteresis is consistent with the interpretation of the irreversible inhibitor layer formation derived from the capacity measurements. The results of measurements carried out at different inhibitor concentrations are shown as Tafel plots in Fig. 5. The Tafel plot for the solution without inhibitor yields an exchange current density of 0.83 mA cmm2 and apparent transfer coefficients of 0.75 and 0.25 for the anodic and cathodic reaction, respectively. These
-2
Cull
N H~Oc..O.l
M Cd’
-3
-6
(’
:
Cull Q.L ” ? s 4 c
L,,.g’
NH~O,,O.lM
CuSO,.p,,=
Oat.AHT
N E u s =
0 LV
-4
G
0.3 ”
0.2 ‘. c ._L 7; 0.1 .’
-5
-6 E ISHEBIV 03
Fig. 4. Current potential curves at the ring-disk electrode in 1 N HZSOI 0.1 M C&O,. Disk: copper; ring: platinum; the disk potential was varied potentiodyuamically with de/dt = 5 mV s-i, rotation frequency = 16.7 s-l, ring potential = 0.4 V for Cu4 detection. cc,,+ = equilibrium concentration of Cu+.
0.2
0.3
0.4
E WiElIV’
Fig. 5. Tafel plot for copper deposition and dissolution at various inhibitor concentrations. For IO-’ M AHT, in the region 0.1 V < E < 0.2 V the cathodic run is shown, otherwise the anodic run is shown.
Inhibition of electrode processes on copper by AHT results agree with the literatureC12, 133. With increasing inhibitor concentration, the exchange current density decreases by 2-3 orders of magnitude and the slopes of the Tafel lines change appreciably. The b factor for the anodic dissolution increases to 125 mV at low overpotentials. The value b = 8 mV at potentials above + 0.47 V does not correspond to a constant surface state since the inhibitor layer dissolves rapidly. On the cathodic side the b factor decreases with increasing inhibitor concentration indicating an increase of the apparent transfer coefficient. For the AHT-free solution we get a sum of the apparent transfer coefficients of almost 1. At the inhibited electrode, on the other hand, LZ+ + a- = 0.7 if both reactions are calculated with z = 2. This assumption, however, may not be true at low anodic overpotentials. Wecan interpret the b factor of 125 mV as due to a dominating reaction (lc), which is too small to be detected at the ring electrode. An alternative interpretation could be based on a shift of the transition state with the changing interfacial structure. The large anodic transfer coefficient of 0.75 can then be explained by the formation of a hydrated Cuz + species involving a transition state several angstroms away from the electrode surface. With increasing inhibitor concentration the thickness of the interfacial layer increases. If the transition state remains at the same distance from the metal surface and is thus within the inhibitor layer, there would be a smaller influence of the interfacial potential drop on the transition state and thus a smaller transfer coefficient. Since the hydration of the transition state in the inhibitor layer will be negligible, the formation of a highly charged Cu2+ ion may be more difficult than that of Cu’ ions. The shift of the relative position of the transition state would cause a simultaneous increase of the transfer coefficient of the cathodic reaction. This indeed is observed, but the shift is not as large as expected due to the change of the dominating process from reaction (la) to reaction (1~). Hence, a decision between these various explanations cannot yet be taken.
827
2H**2e---+H: CullN HfiOl.AHT
-2
-0.8
-0.7
imental results, however, are in complete contradiction of such a model. A probable explanation could be a hydrogen evolution taking place at the surface of the inhibitor layer with the same activation energy as in the case of copper but with a much smaller influence of the
-0.5
-0.L
-0.3
Fig. 6. Tafel plot of cathodic hydrogen evolution on copper without and with inhibitor. Data from potentiodynamic experiments, deaerated solution, &/dt = 5 mV s- ‘.
potential drop since the reaction takes place far away from the metal surface. Cathodic oxygen reduction Since the inhibitor is usually applied in aerated solutions we have to investigate also the cathodic oxygen reduction. Since reactions (4b) and (3) may take place simultaneously with (4a), we first investigated the H202 reduction (3) at the disk electrode. Figure 7
?
Hz02l2 H’+2 e-3 2 t120 r A”,=0\C”11N”2S0~,0.1M HzOz.AHT \
Cathodic hydrogen evolution For many metals cathodic hydrogen evolution is dominant near the corrosion potential. To investigate reaction (2) we measured its current-potential curves in deaerated 1 N H2S04. Without inhibitor, the exchange current density is small (6.6 x lo-’ A cmeZ ) and the b factor is about 120mV corresponding to a cathodic transfer coefficient of 0.5, calculated for the Heyrovsky reaction (Fig. 6). With increasing inhibitor concentration the current density decreases by more than one order of magnitude. The exchange current density remains constant but the Tafel factor increases up to 190 mV corresponding to TV_ = 0.3. If the hydrogen evolution takes place at the metal surface the transfer coefficient should increase as was observed for the cathodic copper deposition. Due to an increasing activation overvoltage i0 should decrease. The exper-
-0.6 c(SHE)IV
k
\
b
I
-0.1
0
0.1 0.2 t(SHE)IV
0.3
0.L
Fig. 7. Tafol plot (eorrccted for diffusion overvoltage) for HxQ reduction on copper with various inhibitor conccntratioas. Rotating copper disk electrode, rotation frequency = 16.7 s-‘. CnIol = 0.01 M. dcjdt PO, = 0 atm, - 5mVs-‘.
J. W. SCHULTZE AND K. WIPPERMANN
828
shows the corresponding Tafel diagram for 0.1 M H202 with the current densities corrected for diffusion overvoltage, using the equation: log idi/&
- i) = log iO - q/b
(3)
where id is the diffusion limited current. The results indicate a fast transfer controlled reaction at the bare copper electrode with i0 = 10-l’ A cm-’ and b = 176 mV indicating a rate determining one-electron transfer reaction. With increasing inhibitor concentration the current density decreases by more than one order of magnitude. However, b remains almost constant. Measurements of the cathodic oxygen reduction (4) at the ring-disk electrode are shown in Fig. 8. At the copper electrode the oxygen reduction takes place as a transfer controlled reaction between 0.2 V and 0 V. At more cathodic potentials diffusion limitation is obvious, and at -0.2 V cathodic hydrogen evolution finally starts. The ring electrode was then kept at 1.2 V to detect the produced H202 and hydrogen. In the range between 0.1 V and 0 V the formation of an increasing amount of H202 is indicated at the ring electrode, but the detectable contribution of reaction (4b) is always small. Below 0 V the H,Oz concentration decreases again, indicating that process (4a) is dominating with only a small contribution from (4b) which is followed by a fast reduction of H202 via
0
-2 c: t _% .z .z
-4
z
DISCUSSION
_.....-‘-
*....-“_
-
-
T’-
.cc -
.)
The infiuence
_iiditf
I
/-
-8
/’
on copper electrodes is a very stable film
I
02+4H’*4e-+ Cu/lNH,SO,,AHT.p~,=la
-1
Cu I1 N HgOt..p,=lat,AHT
r
‘:
2 H&
E u
-2 5 .z z + ..- 1 .= -3 I c”
dc 5 0.4 E -, .z _; =
of concentration
The inhibitor layer formed below 0.4V in acid solution
/ ,
-6
reaction (3). The increase of the ring current below a disk potential of - 0.2 V indicates the participation of cathodic hydrogen evolution. With increasing inhibitor concentration the overall current density at the disk electrode decreases due to the inhibition of the oxygen reduction as well as the hydrogen evolution, and the diffusion limited current id is not reached. Hence, we must conclude that the overall reaction (4a) is strongly inhibited. The ring current shows an inhibition of the Hz02 formation (4b) at E > -0.1 V. At - 0.2 V the contribution of (4b) to the disk current is almost constant, x 0.2 mA cm ‘, but is small by comparison with a disk electrode amount of some mA cm-’ from reaction (4a). Moreover, we have to consider the inhibition of reaction (3) at the Cu surface. If the H202 cannot be reduced there, as was shown above, it will be detected at the ring electrode. Thus a constant ring current of H202 oxidation means in fact a slightly smaller rate of reaction (4b). Since the current of reaction (4b) is negligible the Tafel plot of the overall reaction current density of oxygen reduction shown in Fig. 9 refers to reaction (4a). The current density is again corrected for the diffusion overvoltage. Parallel Tafel lines are obtained for all inhibitor concentrations with a b factor of 139 mV, corresponding to a transfer coefficient of 0.42 calculated for a one-electron transfer reaction. Surprisingly, the inhibition of the oxygen reduction is somewhat smaller than that of HZ02 reduction. This can be explained by the different adsorbabilities of O2 and HZ02 molecules at the AHT surface. The constant Tafel factor for oxygen reduction means that the electron transfer reaction is not influenced by the thickening of the interfacial layer.
0.2
-4 0 -0.3
-0.2
-0.1 0 q,,s< (SHE) I V
Fig. 8. Ring-disk measurements tion at a copper disk. I N H,SO,, IO-’ M, respectively. rotation potential = 1.2 V for
0.1
-0.3 of cathodic oxygen reducPO, = 1 atm, cAHT= 0 and frequency= 16.7 s-l, ring H102 detection.
-0.2
-0.1 E(SHEIIV
0
0.1
0.2
Fig. 9. Tafel plot for cathodic oxygen reductionat Cu. Data from Fig. 8.
Inhibition of electrode processes on copper by AHT
all ion and electron transfer reactions at the interface. Inhibition is observed already in 10e4 M solutions, but increases up to lo-’ M due to the thickening of the inhibitor layer. This is indicated by the decreasing of the capacity down to 2 PFcm ‘. Since we do not know the exact configuration of inhibitor molecules we confine our further discussion to the current density/concentration relationship at constant electrode potential. Figure 10 summarizes the experimental data for all single reactions in a double logarithmic log i/log c plot. At the ordinate the values for pure solutions are given. Nearly straight lines are observed for all reactions, but with differing slopes. The largest slope is observed for the transfer reaction of copper ions which cannot penetrate the inhibitor layer in a fully hydrated state. The smaller oxygen molecules on the other hand may have less difficulty as is well known to be the case for many polymer films. Because of the small influence of the inhibitor concentration, the cathodic hydrogen evolution is assumed to take place at the inhibitor/electrolyte interface, with only a small influence of the potential drop. The true physical reason of the linear log i/log c relationship is not yet clear, but from a formal point of view it can be influencing
analysed in the following way. Similar to the Temkin isothermC13, 143 the coverage and thickness of a species at an interface can be described by a logarithmic relation of the following type:
ii?
:
-5
*Ol”+lo”
-4 lg c*~/mole.
I
-3 I“
Fig. 10. Double logarithmic plot of current density vs concentration for various single reactions of Table 1. Electrode potential is indicated at the lines. Values for cAHT = 0 are indicated at the ordinate.
l/klogc+yFAe+const.
O=d/do=
(4)
with do = thickness of a monolayer, y = electrosorption valency, k = constant. The rate of ion or electron transfer reactions depends on 8 or d often in a similar manner with any fraction a, (es a ‘%hemical transfer coefficient”[13]): log; Combining
= -ma,kl)+;
= --a,&;-+$
(5)
0
(4) and (5) we obtain: (3 log i/dlogc),
= -a,.
(6)
The value of the fraction a, should depend on the type of reaction. For example, we expect that for an electron transfer reaction at an inhibitor layer with metallic conductivity, cl, = 0. But for an ion transfer reaction through a barrier layer, 0 < a, -=z 1. For a reaction with adsorption of the intermediates at a constant surface, eg H or H202 at the inhibitor surface, the value of a, should deviate from 0 only in the case of a limited electronic conductivity. The slope of the lines in Fig. 10 is in qualitative agreement with such a rough description. The infiuence
-6 ..
829
of potential
The results of the kinetic experiments and the Tafel diagrams 5, 6, 7 and 9 are summarized in Table 2. Physical interpretations were already discussed above in corresponding chapters. Figure 11 shows a schematic representation of the interphase with various possible reaction sites at the copper surface and the inhibitor surface. The change of the transfer coefficients parallels roughly the influence of concentration shown in Fig. 10: reaction (1) with a large slope (a log i/a log c)= shows a large shift of the transfer coefficient, while a small (d log i/a log c)~ with an almost constant transfer coefficient is observed for the electron transfer reactions (2 j(4) which could take place at the AHT surface. Interpretation
of rhe overall
current
potential
curves
After these investigations of the single electrode reactions (l)-(4), we can now discuss the overall
Table 2. Results of kinetic measurements in pure solution and 10-j M AHT H+Cl-, respectively. e0 = equilibrium potential in 1 N HZS04, i0 = exchange current density, a + , a_ = apparent transfer coefficients ‘AH,’
(M)
80
P (sW1
i0 (A cm-‘) a+
c
z log i/a log
0 1o-3 0 10-3 0 10-S10)
(Fig.
(1)
Co=CU’+ 0.30 8x 104 2 x 10-e 0.15 0.23 0.25 -0.8 0.44
(2)
+2e-
Hz=2H+
+2e-
(3) ~H~OG=H~O~+ZH+
- 0.021 lx 10-s lx 10-s
1.743 2 x lo-” 6 x lo-l3
0.51 - 0.43 0.32
0.33
+2e-
2H20=02
(W +4H+
1.222 ;;
;;I::
-
-0.56 0.33
0.42 -0.21 0.42
+4e
J. W. SCHULTZEAND K. WIPPERMANN
830
1
N
oxygen reduction and hydrogen evolution with increasing cathodic overvoltage. The broken tines in Fig. 12 indicate the regions of dominating reactions. The free corrosion potential E,,, is 0.2 V in inhibitor free solutions corresponding to a corrosion current density ofabout 12 JIA cm-‘. By addition of lo-’ M inhibitor Emrr shifts to 0.34 V, and due to the inhibition of all single reactions the corrosion current i,, to a smaller current value of 0.5 PA cm- ‘. The shift of E,, can be explained by the well-established fact that the inhibitor has a stronger influence on the copper dissolution reaction than on the cathodic oxygen reduction. That means that AHT has to be classified as an inhibitor which acts perferably on the anodic side. For practical application this is no problem since the inhibitor layer is stable up to 0.45 V which is far beyond the limit reached in real systems.
H2S04
Acknowledgements-The authors gratefully acknowledge the fruitful discussions with Dr Schnegelberger, Dr Penninger and Dr PreuR (Firma Henkel KGaA Dusseldorf). We are obliged to Professor Strehblow and H. D. Speckmann for lending the ring-disc electrode equipment. Fig. 11. Schematic representation of the interphase, showing
the different reaction sites at the copper surface and the inhibitor surface, respectively.
REFERENCES
current-potential curves recorded in aerated HzSO_, solutions. Figure 12 shows the Tafel plot. To simulate the practical conditions we recorded the currentpotential curves after small amounts of Cu corrosion which yields, of course, a less-defined copper concentration before the metal surface. We estimate c,-~z+ < lo-“ M and i_,,, < lo-“ A cm-*. By comparison with the Tafel plots of the single reactions we can interpret the current-potential curves by copper dissolution on the anodic side and by copper deposition,
:: 3. 4. 5. 6. 7. 8.
Procter & Gamble, Br. Pat. 652 339, Dec. (1947). J. B. Cotton and I. R. Scholes, Br. Corros. J. 2, 1 (1967). G. W. Poling, Corros. Sci. 10,359 (1970). Henkel KGaA, Offenlegungsschrift DE2934461Al (1981). J. Penninger, unpublished results. K. Wippermann, Diploma&sit, Universitiit Dusseldorf (1986). K. Wippermann, J. Penninger and J. W. Schultze, in preparation (1987). M. M. Lohrengel, K. Schubert and J. W. Schultze, Werksr. Korros. 32, 13 (1981).
I
I
-1.0
‘t
‘., Cu+GJ2’
pJ2+~ cu ,
I -0.5
0 E
0.5
(SHEIIV
Fig. 12. Tafel plot for the overall current potential curves on copper electrodes in 1 N H2S04 with various inhibitor concentrations. PO* = 0.2 atm, ds/dr = 5 mV s- ‘. The vertical broken lines separate regions of dominating reactions.
Inhibition of electrode processes on copper by AHT 9. H. D. Speckmann, M. M. Lohrengel, J. W. Schultze and H.-H. Strehblow, Ber. Bunsenws. phys. Chem. 89, 392 (1985). 10. B. P. L&he1 and H.-H. Strehblow, Werkst. Korros. 31, 353 (1980).
831
11. A. Frumkin, Z. Phys. 35, 791 (1926). 12. J.O’M. Bockris and M. Enyo, Trans. Faraday Sot. 58, 1187 (1962). 13. J. W. Schultze, Ber. Bunsenges. phys. Chem. 74,705 (1970). 14. M. Temkin, J. phys. Chem. USSR 15, 296 (1941).
OF ELECTRODE AHT IN ACID J. W.SCHULTZE
PROCESSES SOLUTIONS
ON
COPPER
BY
and K. WIPPERMANN
Institut fur Physikalisohe Chemie II der Universitit Diisseldorf, UniversitHtsstraBe 1, D-4000 Dusseldorf, F.R.G. (Received 21 August
1986; in revised firm
17 November 1986)
Aktract-3-Atnino-5-heptyl-1,2,4-triazole (AHT)* is a very effective inhibitor of copper corrosion in neutral as well as acid solutions. From capacity-potential curves and their time dependence it follows that a compound layer of AHT (Cu, Cl) of 1-2 nm thickness is formed which is stable up to 0.4 V. At E > 0.45 V AHT is dissolved immediately. Current-potential curves were measured for all possible electrode reactions. The inhibitor affects the anodic oxidation of copper strongly. The cathodic oxygen reduction is influenced as well but to a lesser degree. The inhibition of the cathodic hydrogen evolution is only important in deaerated solutions and at large cathodic overvoltages. From the values and the shift of transfer coefficients we must conclude that the reaction site differs for transfer reactions of electrons, protons and copper ions. The influence of the AHT concentration is exnlained by the thickening of the layer, the conductivity of which differs for ions and electrons.
INTRODUCTION
addition of 10m3 M AHT causes’s decrease of the current density to very small values in the potential range - 0.5 to + 0.4 V. The capacity-potential curves show that this inhibition efficiency is combined with a strong decrease of the electrode capacity by a factor of almost l/10. At very high potentials only the increase of the capacity and the anodic current density indicates a desorption of the inhibitor. To investigate the electrode reactions (l)-(5) we chose corresponding systems, both without (cAHT = 0) and with inhibitor (10m5 M < cAHT 6 lo-’ M), as follows:
It has been well known for a long time that benzotriazole inhibits the copper corrosion in neutral and alkaline solutions, but it is less effective in acid solutions[lG3]. The inhibition is due to the formation of thick layers of copper-benzotriazole complexes and probably also due to the formation of copper oxide layers protected by the inhibitor. A few years ago the inhibitor 3-amino-S-alkyl-1,2,4-triazole (Trademark Texamin AT 1) was shown to have the same or even better inhibition efficiency especially in acid -Alutions[4--6]. In spite of the practical success of this inhibitor no scientific investigations of the manifold reactions which can take place in this system in neutral solutions have been carried out. Table 1 shows a summary of possible reactions (l)-(S) and the corresponding electrode potentials. Preliminary investigations in neutral solutions showed a complex overlap of anodic corrosion, oxide formation and inhibitor layer formation.
System (1) 1 N HISO+, PO, = 0 atm, 0.1 M Cu” (2) 1 N H2S04, PO, = 0 atm (3) 1 N HZS04, PO, = 0 atm, 0.1 M H202 (4) 1 N HZS04, PO, = 0.2 and 1 atm
Electrode reaction (see Table 1) (lak(W (2) (3)
Pa) + (4b)
The results of the measurements will be presented below. Further results regarding the influence of pH, anions, temperature, the comparison with benzotriazole and XPS measurements will be presented in Ref. [7].
Scheme 1. To simplify the system we carried out extended investigations in acid solutions where we could exclude the oxide formation reactions [(ld)-(lf)]. Figure 1 shows examples of current-potential and capacity-potential curves in 1 N HZSO, both with and without inhibitor. The anodic current density is due to the dissolution of copper (la), while the cathodic reaction at small overvoltages is caused by oxygen reduction (4) or at large overvoltages by hydrogen evolution (2). The
EXPERIMENTAL The electronic set-up of the electrochemical measurements is described elsewhere[S]. It consisted of a fast rise potentiostat, an X-Y recorder and a signal source and lock-in amplifier for the UC measurements. Chemicals
*Commercial inhibitor of Henkel KGaA with Trademark Texamin AT 1.
1 N sulfuric acid with a pH of 0.4 was used as electrolyte. It was prepared with p.A. HZS04 (Merck)
823
824
J. W. SCHULTZE AND K. WIPPERMANN
Table
1. Possible reactions at copper electrodes in aqueous solutions Standard potential Reactions
CWWI
(la) Cu +cu=++2e(lb) Cu *cu++e(lc) cu +*cw++.(Id) 2Cu+H~O=Cu,O+2H++-te(le) CuZO + HZ0 =2Cu0+2H++2em (If) Cu,O+ 3HZO&2Cu(OH)Z +2H+ (2) H,=2H+ +2e(3) 2HzO &H202+2Hf+2e(4a) 2H20=0, +4H+ +4e(4b) H20,==0,+2H+ +2e(5) AATw + Cu G= (Cu-AAT)+ (5b) AAT + Cu + F= (Cu-AAT) (5~) AAT + Cu2+ + (Cu-AAT)‘+
0.4
‘: E Y $r T=
Cull N H~O&,;0.2
+2e-
+ 0.336 + 0.520 +0.153 +0.471 + 0.669 + 0.747 0.000 + 1.770 + 1.229 +0.682 -
at.AHT
Cell and electrodes The electrochemical measurements were carried out in a vessel of duran glass thermostated at 25°C. As working electrode we used copper cathodically deposited on a gold wire[9]. The gold wire had a diameter of 0.4 mm and an area of 0.25 cm’. A gold foil of 4 cm2 was used as counter electrode. To check the influence of the crystal plane we used a (11 l)-oriented copper single crystal, which was kindly supplied by Professor Schaarwachter (Universitat Dortmund). The single crystal plane was mechanically polished using diamond spray (15 p, 3 p) and finally electropolished in 75% H3P04 with an anodic current density of 0.4 A cm-*. For the detection of intermediates we used a rotating split ring-disc electrode which consisted of a central copper disc (99.999 o/0Cu) of 3.4 mm diameter surrounded by 2 Pt half-rings[lO]. The collection efficiency of 1 Pt half-ring was determined to be N = 0.15 in agreement with calculations from geometry. The rrd electrode was mechanically polished with 3 p diamond spray and washed with alcohol and water before each experiment. RESULTS
0.2 0
Capacity
measurements
Fig. 1. (a) Potentiodynamic current potential curves recorded at a copper (111) single crystal in 1 N H2S04 with cAHT = 10e3 M 1-l and 0 M I- 1 respectively. Aerated solution, d&/d/dr = 5 mV s- ‘. (b) Electrode capacity under the same conditions. 1 kHz, 2 mV amplitude.
Since the electrode capacity indicates structural changes at the interface, we recorded the capacity in inhibitor solutions of various concentrations. 1N H,SO, without Cu2+ was chosen so that we could extend the measurements over a large range of electrode potentials. Systems 2 and 4 (deaerated and aerated) gave equal results. Hence we always used system 4. Further, there was no difference between single crystal surfaces and the polycrystalline Cu surface. Hence, polycrystalline material was always used. Figure 2 shows the electrode capacity measured during a slow potentiodynamic sweep under stabilized conditions (after more than 5 cycles). In 10m4 M inhibitor solution one observes already a strong decrease of electrode capacity which then becomes very small in 10e3 M and lo-* M solutions. Capacity values of indicate the formation of a very about S PF cm-’ stable inhibitor layer in the entire accessible potential range. Using the Frumkin model of parallel condensors for the adsorption of organic molecules, we estimated the coverage of AHT, 13,using the equation (1): @= (c,=.-c)/c,=.-c*~,). (1)
and three times distilled water. The solutions were either aerated by oxygen (99.999%) or de-aerated by nitrogen (99.999 %)_ The inhibitor 3-amino-5-heptyl1,2,4-triazole (AHT) was used as a hydrochloride which has a better inhibition efficiency than its sulfate form[7]. The p& value of AHT was determined as 9.35. The saturation concentration of AHT hydrochloride in water is 0.1 mol l- ‘. Correspondingly the standard free enthalpy of the solution 15.6 kJ mol- ‘. The inhibitor was kindly supplied by Henkel KGaA Dusseldorf.
Taking C,= r to be the minimum value of C for a lo-’ M AHT solution, Ca= r = Cmin = 2.5 pF cme2, we calculate (3 > 0.9 for lo- 3 M solutions and 8 z 1 for lo-’ M AHT. Such an estimation, however, is misleading, since AHT is not adsorbed reversibly under the conditions of Fig. 2, which correspond to an almost steady state of the interface. Figure 3 shows that for a fresh surface (curve 1) the capacity decreases continuously during cyclization if the desorption potential of 0.45 V is not exceeded. An almost constant value of about 5 PF cm -2 is reached in the 5th cycle. Exceeding 0.45 V, the inhibitor dissolves and the capacity increases up to 60 PF cm-’ and remains high during the cathodic cycle Sb.
-0.2 50
40 v $ 30 Y a. z 20
10
0 -1.0
0
-0.5
0.5
E(SHEIIV
Inhibition of electrode processes on copper by AHT
CullN
H2SOI,pO;0.2 at,AHT
/ /
I ; cAHT=O R-\
,
\
I
\
c-’
\
w--e
a
-0.4
I
\
,
.
825
\
-0:z
0
2’
0:2
0.4
E(SHE)IV
Fig. 2. Electrode capacity of copper in 1 N H,SO, with various concentrations of inhibitor under steady state conditions. Aerated solution &/dt = 5 mV s-l_
80
Cull
NH,S0,,p,2=
0.2 at,lO-‘M
AHT
60
0
-0.4
-0.2
0
0.2
0.4
flSHE)/V
Fig. 3. lst, 2nd and 5th potentioclynamic capacity-potential curves in 1 N H2SO4 after addition of lo-’ hi AHT. Broken line gives results of potentiostatic pulse measurements from + 0.5 V (free surfarr) to 0.3,O and - 0.5 V, respectively. To test the reversibility of the adsorption system we carried out additional potentiostatic step experiments, starting from a very positive potential, 0.5 V, where the inhibitor is des0rbe.d completely. Then we stepped the potential to the measuring potential and recorded the electrode capacity time-dependence. At f0.3 V the steady state value of Fig. 2 is approached in a short time but at lower potentials we observed a very slow decrease of the electrode capacity indicating the slow formation of the inhibitor layer. Since molecular adsorption reactions are usually fast and diffusion
limitation is to be expected for a few seconds only, an adsorption equilibrium should be maintained in few seconds. Hence, the strong deviation between the potential step experiments and the curve 1 in Fig. 3 on the one hand, and the steady state potentiodynamic sweep experiments in Fig. 2 and curve 5a in Fig. 3 on the other hand, indicate that the inhibitor layer formed in the steady state is not an adsorption layer but a stable precipitate which is formed slowly and irreversibly and does not change with electrode potential. Finally, to compare these results with a simple
826
J. W. SCHULTZEAND
adsorption system of AHT we measured the electrode capacity on mercury electrodes[GJ. For this system the lowest capacity values observed were about clearly exceeding the values observed on 8 PFcm-‘, copper electrodes. Hence, we conclude again that the AHT layer on copper electrodes cannot be described by a simple adsorption system (5a) of Table 1 but rather by equations (Sb) and (5c), respectively. More evidence for this interpretation was obtained from electrolyte exchange experiments and XPS measurements after extended rinsing of the electrode with pure water[7]. From the XPS measurements we conclude that the thickness of the inhibitor layer at the free corrosion potential is within the range of 1-2 nm. Taking the condenser formula d = DDo /C
(2)
we obtain from C = 2.5 FF cm-‘, a dielectric constant of D = 5, a film thickness of 1.8 nm, which is a reasonable value for a layer of inhibitor molecules. Measurements in sulphuric acid with AHT sulfate[7] and AHT chloride yielded smaller capacities for the chloride solution indicating a chloride catalysed formation of the inhibitor layer. Hence, we must conclude that in the system Cu/10m3 M AHT chloride, layers of 1-2 molecules are formed in the steady state, containing an unknown amount of cosorbed copper and chloride ions. Therefore, the influence of adsorption is small or negligible, and it is clear that polycrystalline electrodes and Conner single crvstals behave equally in AHT solution. _-
K. WIPPERMANN Copper
dissolution
and deposition
We carried out investigations of the anodic copper dissolution and cathodic copper deposition in system 1 using a rotating ring-disk electrode in 0.1 M Cur+. The upper part of Fig. 4 demonstrates the dependence of the disk electrode current density on the electrode potential. In solutions without inhibitor we see the transfer controlled reaction with large transfer currents as seen in Fig. 4a between 0.2 V and 0.35 V. The lower part of Fig. 4 shows the corresponding ring current at a potential of0.4 V, indicating the formation of Cu+. The ring current density increases with increasing potential almost parallel to the equilibrium concentration of Cu + calculated from the Nernst equation for reaction (1 b). In solutions with 1O-3 M AHT the current density at the disk as well as at the ring is almost 0 in the range from 0.2 V to 0.4 V (Fig. 4a). At E > 0.47 V the rapid rise in the disk current indicates desorption of the inhibitor which allows anodic dissolution of copper. This current decreased more slowly for the cathodic sweep. Simultaneously Cu+ is indicated at the ring electrode. Below 0.2 V a small amount of copper deposition takes place with a current density of about without production of Cu+. The slow 1 mA cm-‘, hysteresis observed for the cathodic current can be attributed to the production of a fresh surface during the copper deposition which allows a faster transfer of copper ions in the following anodic sweep, up to 0.2 V. The existence of this hysteresis is consistent with the interpretation of the irreversible inhibitor layer formation derived from the capacity measurements. The results of measurements carried out at different inhibitor concentrations are shown as Tafel plots in Fig. 5. The Tafel plot for the solution without inhibitor yields an exchange current density of 0.83 mA cmm2 and apparent transfer coefficients of 0.75 and 0.25 for the anodic and cathodic reaction, respectively. These
-2
Cull
N H~Oc..O.l
M Cd’
-3
-6
(’
:
Cull Q.L ” ? s 4 c
L,,.g’
NH~O,,O.lM
CuSO,.p,,=
Oat.AHT
N E u s =
0 LV
-4
G
0.3 ”
0.2 ‘. c ._L 7; 0.1 .’
-5
-6 E ISHEBIV 03
Fig. 4. Current potential curves at the ring-disk electrode in 1 N HZSOI 0.1 M C&O,. Disk: copper; ring: platinum; the disk potential was varied potentiodyuamically with de/dt = 5 mV s-i, rotation frequency = 16.7 s-l, ring potential = 0.4 V for Cu4 detection. cc,,+ = equilibrium concentration of Cu+.
0.2
0.3
0.4
E WiElIV’
Fig. 5. Tafel plot for copper deposition and dissolution at various inhibitor concentrations. For IO-’ M AHT, in the region 0.1 V < E < 0.2 V the cathodic run is shown, otherwise the anodic run is shown.
Inhibition of electrode processes on copper by AHT results agree with the literatureC12, 133. With increasing inhibitor concentration, the exchange current density decreases by 2-3 orders of magnitude and the slopes of the Tafel lines change appreciably. The b factor for the anodic dissolution increases to 125 mV at low overpotentials. The value b = 8 mV at potentials above + 0.47 V does not correspond to a constant surface state since the inhibitor layer dissolves rapidly. On the cathodic side the b factor decreases with increasing inhibitor concentration indicating an increase of the apparent transfer coefficient. For the AHT-free solution we get a sum of the apparent transfer coefficients of almost 1. At the inhibited electrode, on the other hand, LZ+ + a- = 0.7 if both reactions are calculated with z = 2. This assumption, however, may not be true at low anodic overpotentials. Wecan interpret the b factor of 125 mV as due to a dominating reaction (lc), which is too small to be detected at the ring electrode. An alternative interpretation could be based on a shift of the transition state with the changing interfacial structure. The large anodic transfer coefficient of 0.75 can then be explained by the formation of a hydrated Cuz + species involving a transition state several angstroms away from the electrode surface. With increasing inhibitor concentration the thickness of the interfacial layer increases. If the transition state remains at the same distance from the metal surface and is thus within the inhibitor layer, there would be a smaller influence of the interfacial potential drop on the transition state and thus a smaller transfer coefficient. Since the hydration of the transition state in the inhibitor layer will be negligible, the formation of a highly charged Cu2+ ion may be more difficult than that of Cu’ ions. The shift of the relative position of the transition state would cause a simultaneous increase of the transfer coefficient of the cathodic reaction. This indeed is observed, but the shift is not as large as expected due to the change of the dominating process from reaction (la) to reaction (1~). Hence, a decision between these various explanations cannot yet be taken.
827
2H**2e---+H: CullN HfiOl.AHT
-2
-0.8
-0.7
imental results, however, are in complete contradiction of such a model. A probable explanation could be a hydrogen evolution taking place at the surface of the inhibitor layer with the same activation energy as in the case of copper but with a much smaller influence of the
-0.5
-0.L
-0.3
Fig. 6. Tafel plot of cathodic hydrogen evolution on copper without and with inhibitor. Data from potentiodynamic experiments, deaerated solution, &/dt = 5 mV s- ‘.
potential drop since the reaction takes place far away from the metal surface. Cathodic oxygen reduction Since the inhibitor is usually applied in aerated solutions we have to investigate also the cathodic oxygen reduction. Since reactions (4b) and (3) may take place simultaneously with (4a), we first investigated the H202 reduction (3) at the disk electrode. Figure 7
?
Hz02l2 H’+2 e-3 2 t120 r A”,=0\C”11N”2S0~,0.1M HzOz.AHT \
Cathodic hydrogen evolution For many metals cathodic hydrogen evolution is dominant near the corrosion potential. To investigate reaction (2) we measured its current-potential curves in deaerated 1 N H2S04. Without inhibitor, the exchange current density is small (6.6 x lo-’ A cmeZ ) and the b factor is about 120mV corresponding to a cathodic transfer coefficient of 0.5, calculated for the Heyrovsky reaction (Fig. 6). With increasing inhibitor concentration the current density decreases by more than one order of magnitude. The exchange current density remains constant but the Tafel factor increases up to 190 mV corresponding to TV_ = 0.3. If the hydrogen evolution takes place at the metal surface the transfer coefficient should increase as was observed for the cathodic copper deposition. Due to an increasing activation overvoltage i0 should decrease. The exper-
-0.6 c(SHE)IV
k
\
b
I
-0.1
0
0.1 0.2 t(SHE)IV
0.3
0.L
Fig. 7. Tafol plot (eorrccted for diffusion overvoltage) for HxQ reduction on copper with various inhibitor conccntratioas. Rotating copper disk electrode, rotation frequency = 16.7 s-‘. CnIol = 0.01 M. dcjdt PO, = 0 atm, - 5mVs-‘.
J. W. SCHULTZE AND K. WIPPERMANN
828
shows the corresponding Tafel diagram for 0.1 M H202 with the current densities corrected for diffusion overvoltage, using the equation: log idi/&
- i) = log iO - q/b
(3)
where id is the diffusion limited current. The results indicate a fast transfer controlled reaction at the bare copper electrode with i0 = 10-l’ A cm-’ and b = 176 mV indicating a rate determining one-electron transfer reaction. With increasing inhibitor concentration the current density decreases by more than one order of magnitude. However, b remains almost constant. Measurements of the cathodic oxygen reduction (4) at the ring-disk electrode are shown in Fig. 8. At the copper electrode the oxygen reduction takes place as a transfer controlled reaction between 0.2 V and 0 V. At more cathodic potentials diffusion limitation is obvious, and at -0.2 V cathodic hydrogen evolution finally starts. The ring electrode was then kept at 1.2 V to detect the produced H202 and hydrogen. In the range between 0.1 V and 0 V the formation of an increasing amount of H202 is indicated at the ring electrode, but the detectable contribution of reaction (4b) is always small. Below 0 V the H,Oz concentration decreases again, indicating that process (4a) is dominating with only a small contribution from (4b) which is followed by a fast reduction of H202 via
0
-2 c: t _% .z .z
-4
z
DISCUSSION
_.....-‘-
*....-“_
-
-
T’-
.cc -
.)
The infiuence
_iiditf
I
/-
-8
/’
on copper electrodes is a very stable film
I
02+4H’*4e-+ Cu/lNH,SO,,AHT.p~,=la
-1
Cu I1 N HgOt..p,=lat,AHT
r
‘:
2 H&
E u
-2 5 .z z + ..- 1 .= -3 I c”
dc 5 0.4 E -, .z _; =
of concentration
The inhibitor layer formed below 0.4V in acid solution
/ ,
-6
reaction (3). The increase of the ring current below a disk potential of - 0.2 V indicates the participation of cathodic hydrogen evolution. With increasing inhibitor concentration the overall current density at the disk electrode decreases due to the inhibition of the oxygen reduction as well as the hydrogen evolution, and the diffusion limited current id is not reached. Hence, we must conclude that the overall reaction (4a) is strongly inhibited. The ring current shows an inhibition of the Hz02 formation (4b) at E > -0.1 V. At - 0.2 V the contribution of (4b) to the disk current is almost constant, x 0.2 mA cm ‘, but is small by comparison with a disk electrode amount of some mA cm-’ from reaction (4a). Moreover, we have to consider the inhibition of reaction (3) at the Cu surface. If the H202 cannot be reduced there, as was shown above, it will be detected at the ring electrode. Thus a constant ring current of H202 oxidation means in fact a slightly smaller rate of reaction (4b). Since the current of reaction (4b) is negligible the Tafel plot of the overall reaction current density of oxygen reduction shown in Fig. 9 refers to reaction (4a). The current density is again corrected for the diffusion overvoltage. Parallel Tafel lines are obtained for all inhibitor concentrations with a b factor of 139 mV, corresponding to a transfer coefficient of 0.42 calculated for a one-electron transfer reaction. Surprisingly, the inhibition of the oxygen reduction is somewhat smaller than that of HZ02 reduction. This can be explained by the different adsorbabilities of O2 and HZ02 molecules at the AHT surface. The constant Tafel factor for oxygen reduction means that the electron transfer reaction is not influenced by the thickening of the interfacial layer.
0.2
-4 0 -0.3
-0.2
-0.1 0 q,,s< (SHE) I V
Fig. 8. Ring-disk measurements tion at a copper disk. I N H,SO,, IO-’ M, respectively. rotation potential = 1.2 V for
0.1
-0.3 of cathodic oxygen reducPO, = 1 atm, cAHT= 0 and frequency= 16.7 s-l, ring H102 detection.
-0.2
-0.1 E(SHEIIV
0
0.1
0.2
Fig. 9. Tafel plot for cathodic oxygen reductionat Cu. Data from Fig. 8.
Inhibition of electrode processes on copper by AHT
all ion and electron transfer reactions at the interface. Inhibition is observed already in 10e4 M solutions, but increases up to lo-’ M due to the thickening of the inhibitor layer. This is indicated by the decreasing of the capacity down to 2 PFcm ‘. Since we do not know the exact configuration of inhibitor molecules we confine our further discussion to the current density/concentration relationship at constant electrode potential. Figure 10 summarizes the experimental data for all single reactions in a double logarithmic log i/log c plot. At the ordinate the values for pure solutions are given. Nearly straight lines are observed for all reactions, but with differing slopes. The largest slope is observed for the transfer reaction of copper ions which cannot penetrate the inhibitor layer in a fully hydrated state. The smaller oxygen molecules on the other hand may have less difficulty as is well known to be the case for many polymer films. Because of the small influence of the inhibitor concentration, the cathodic hydrogen evolution is assumed to take place at the inhibitor/electrolyte interface, with only a small influence of the potential drop. The true physical reason of the linear log i/log c relationship is not yet clear, but from a formal point of view it can be influencing
analysed in the following way. Similar to the Temkin isothermC13, 143 the coverage and thickness of a species at an interface can be described by a logarithmic relation of the following type:
ii?
:
-5
*Ol”+lo”
-4 lg c*~/mole.
I
-3 I“
Fig. 10. Double logarithmic plot of current density vs concentration for various single reactions of Table 1. Electrode potential is indicated at the lines. Values for cAHT = 0 are indicated at the ordinate.
l/klogc+yFAe+const.
O=d/do=
(4)
with do = thickness of a monolayer, y = electrosorption valency, k = constant. The rate of ion or electron transfer reactions depends on 8 or d often in a similar manner with any fraction a, (es a ‘%hemical transfer coefficient”[13]): log; Combining
= -ma,kl)+;
= --a,&;-+$
(5)
0
(4) and (5) we obtain: (3 log i/dlogc),
= -a,.
(6)
The value of the fraction a, should depend on the type of reaction. For example, we expect that for an electron transfer reaction at an inhibitor layer with metallic conductivity, cl, = 0. But for an ion transfer reaction through a barrier layer, 0 < a, -=z 1. For a reaction with adsorption of the intermediates at a constant surface, eg H or H202 at the inhibitor surface, the value of a, should deviate from 0 only in the case of a limited electronic conductivity. The slope of the lines in Fig. 10 is in qualitative agreement with such a rough description. The infiuence
-6 ..
829
of potential
The results of the kinetic experiments and the Tafel diagrams 5, 6, 7 and 9 are summarized in Table 2. Physical interpretations were already discussed above in corresponding chapters. Figure 11 shows a schematic representation of the interphase with various possible reaction sites at the copper surface and the inhibitor surface. The change of the transfer coefficients parallels roughly the influence of concentration shown in Fig. 10: reaction (1) with a large slope (a log i/a log c)= shows a large shift of the transfer coefficient, while a small (d log i/a log c)~ with an almost constant transfer coefficient is observed for the electron transfer reactions (2 j(4) which could take place at the AHT surface. Interpretation
of rhe overall
current
potential
curves
After these investigations of the single electrode reactions (l)-(4), we can now discuss the overall
Table 2. Results of kinetic measurements in pure solution and 10-j M AHT H+Cl-, respectively. e0 = equilibrium potential in 1 N HZS04, i0 = exchange current density, a + , a_ = apparent transfer coefficients ‘AH,’
(M)
80
P (sW1
i0 (A cm-‘) a+
c
z log i/a log
0 1o-3 0 10-3 0 10-S10)
(Fig.
(1)
Co=CU’+ 0.30 8x 104 2 x 10-e 0.15 0.23 0.25 -0.8 0.44
(2)
+2e-
Hz=2H+
+2e-
(3) ~H~OG=H~O~+ZH+
- 0.021 lx 10-s lx 10-s
1.743 2 x lo-” 6 x lo-l3
0.51 - 0.43 0.32
0.33
+2e-
2H20=02
(W +4H+
1.222 ;;
;;I::
-
-0.56 0.33
0.42 -0.21 0.42
+4e
J. W. SCHULTZEAND K. WIPPERMANN
830
1
N
oxygen reduction and hydrogen evolution with increasing cathodic overvoltage. The broken tines in Fig. 12 indicate the regions of dominating reactions. The free corrosion potential E,,, is 0.2 V in inhibitor free solutions corresponding to a corrosion current density ofabout 12 JIA cm-‘. By addition of lo-’ M inhibitor Emrr shifts to 0.34 V, and due to the inhibition of all single reactions the corrosion current i,, to a smaller current value of 0.5 PA cm- ‘. The shift of E,, can be explained by the well-established fact that the inhibitor has a stronger influence on the copper dissolution reaction than on the cathodic oxygen reduction. That means that AHT has to be classified as an inhibitor which acts perferably on the anodic side. For practical application this is no problem since the inhibitor layer is stable up to 0.45 V which is far beyond the limit reached in real systems.
H2S04
Acknowledgements-The authors gratefully acknowledge the fruitful discussions with Dr Schnegelberger, Dr Penninger and Dr PreuR (Firma Henkel KGaA Dusseldorf). We are obliged to Professor Strehblow and H. D. Speckmann for lending the ring-disc electrode equipment. Fig. 11. Schematic representation of the interphase, showing
the different reaction sites at the copper surface and the inhibitor surface, respectively.
REFERENCES
current-potential curves recorded in aerated HzSO_, solutions. Figure 12 shows the Tafel plot. To simulate the practical conditions we recorded the currentpotential curves after small amounts of Cu corrosion which yields, of course, a less-defined copper concentration before the metal surface. We estimate c,-~z+ < lo-“ M and i_,,, < lo-“ A cm-*. By comparison with the Tafel plots of the single reactions we can interpret the current-potential curves by copper dissolution on the anodic side and by copper deposition,
:: 3. 4. 5. 6. 7. 8.
Procter & Gamble, Br. Pat. 652 339, Dec. (1947). J. B. Cotton and I. R. Scholes, Br. Corros. J. 2, 1 (1967). G. W. Poling, Corros. Sci. 10,359 (1970). Henkel KGaA, Offenlegungsschrift DE2934461Al (1981). J. Penninger, unpublished results. K. Wippermann, Diploma&sit, Universitiit Dusseldorf (1986). K. Wippermann, J. Penninger and J. W. Schultze, in preparation (1987). M. M. Lohrengel, K. Schubert and J. W. Schultze, Werksr. Korros. 32, 13 (1981).
I
I
-1.0
‘t
‘., Cu+GJ2’
pJ2+~ cu ,
I -0.5
0 E
0.5
(SHEIIV
Fig. 12. Tafel plot for the overall current potential curves on copper electrodes in 1 N H2S04 with various inhibitor concentrations. PO* = 0.2 atm, ds/dr = 5 mV s- ‘. The vertical broken lines separate regions of dominating reactions.
Inhibition of electrode processes on copper by AHT 9. H. D. Speckmann, M. M. Lohrengel, J. W. Schultze and H.-H. Strehblow, Ber. Bunsenws. phys. Chem. 89, 392 (1985). 10. B. P. L&he1 and H.-H. Strehblow, Werkst. Korros. 31, 353 (1980).
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11. A. Frumkin, Z. Phys. 35, 791 (1926). 12. J.O’M. Bockris and M. Enyo, Trans. Faraday Sot. 58, 1187 (1962). 13. J. W. Schultze, Ber. Bunsenges. phys. Chem. 74,705 (1970). 14. M. Temkin, J. phys. Chem. USSR 15, 296 (1941).