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Considerations on the mechanism of the hydrogenation of organic compounds in aqueous solutions on noble metal catalysts
Bktrochimica
Acta. 1970. Vol. 15, pp. 987 to 997.
Pcrktunon Press. Printed in Northern Ireland
CONSIDERATIONS ON THE MECHANISM OF THE HYDROGENATION OF ORGANIC COMPOUNDS IN AQUEOUS SOLUTIONS ON NOBLE METAL CATALYSTS* C. WAGNER Max-Planck-lnstitut fllr physikahsche Chemie, GGttingen, Bundesrepublik, Deutschland Abstract-Hydrogenation of an organic compound A on Pt as catalyst in aqueous solutions may occur (1) by a sequence of non-electiochemical reactions involving adsorbed hydrogen atoms and radicals as intermediates and (2) by anodic ionization of hydrogen and consecutive cathodic reduction of compound A. This paper is concerned with experimental methods that are especially promising for deciding whether a non-electrochemical or an electrochemical mechanismprevails for a particular compound. R&um&L’hydrog&ration d’un compose organique A en solutions aqueuses, avec le Pt comme catalyseur, peut se manifester comme une succession de reactions non electrochimiqu~, impliquant comme intermecliaires,Padsorption d’atomes d’hydro@ne et de radicaux, l’ionisation anodique de l’hydrogene et la reduction cathodique consecutive du compose A. Le present memoire disc&e ies methodes experimentales les plus inttkessantes,en vue de decider si urt type de m&anisme, non6lectrochimique ou &ctrochimique, prevaut pour un cornpod particulier. Zussammenfassrmg-Die Hydrierung einer organischen Verbindung A an Pt als Katalysator karm entweder durch eine Folge nichtelektrochemischer Reaktionen mit adsorbierten Wasserstoffatomen .und Radikalen als Zwischenprodukt oder durch anodische Ionisierung von WasserstofI und kathodische Reduktion der Verbmdung A erfolgen. In dieser Arbeit werden experimentelle Methoden angegeben, die besonders gee&net scheinen, um zu entscheiden, ob em nichtelektrochemischer oder ein elektrochemischer Mechanismus im Einzelfall vorliegt. STATEMENT
OF THB
PROBLEM
THE HYDROGENATION of an organic compound A by H, to AH, in an aqueous solution
with the help of a noble metal, eg platinum, as a catalyst, may occur via various mechanisms.lve First, hydrogen molecules may dissociate on the surface of the noble metal and subsequently hydrogen atoms may react with compound A in consecutive steps in the same way as the catalytic hydrogenation of a gaseous organic compound, eg ethylene at a metal/gas interface proceeds with adsorbed hydrogen atoms and adsorbed AH radicals as intermediates (mechanism I) Hs(aq) H,(ad) A(aq) -
Ha(ad), 2 II(
(2)
A(ad),
(3)
A(ad) + H(ad) - AH(ad), AH(ad) + H(ad) - AH,(ad),
(1)
(4) (5)
AH&d) - AH&q). (6) Secondly, electrochemical mechanisms involving consecutive anodic and cathodic reactions with zero net current at the metal/solution interface may be considered. Investigations on the cathodic reduction of quinone to hydroquinone by Vetterf have shown that different electron-transfer reactions prevail at low pH values (pH < 5) and high pH values (pH > 6). In view of these results one may consider the following sequences of reactions with prevalence of reactions on the left-hand side (mechanism IIa) at low pII values and prevalence of reactions on the right-hand side * Manuscript received 10 April 1969. 987
C. WAGNER
988
{mechanism
IIb) at high pH values, H,(aq) -+ H,(ad), H,(ad) - 2 H(ad), A(aq) - A(ad), 2(H(ad) -+ H+(aq)
A(ad) + H+(aq) ---f AH+(ad), AH+(acl) + e- AH(ad)
+ H+(aq) -
AH$(ad)
+ e- -
(1) (2) (3) + e-1;
AH(ad),
A-(ad)
AHz+(ad),
(9) (10)
AH,(ad) ,
(11)
AH-(ad)
AH,(ad)
(7)
A(ad) + e- -
(8)
+ H+(aq) -
A-(ad),
(12)
AH(ad),
+ e- -
AH-(ad),
(13) (14)
+ H+(aq) -
AH,(ad),
(15)
AH(ad) -+ AH,(aq).
(6) One may hypothesize that mechanism I prevails in the case of hydrogenation of unsaturated hydrocarbons and conversely electrochemical mechanism (IIa) or (IIb) prevails in systems where the equilibrium potential of the redox couple A-AH, at a noble metal is readily established, ie the exchange cd of the couple A-AH, is relatively high, as in the case of the quinone-hydroquinone couple. It is the objective of this paper to outline experimental investigations which are promising in order to decide which mechanism is instrumental for hydrogenation of a particular inorganic or organic compound A. SYMBOLS
We use the following: E, electrode potential of a platinum .foil in a solution containing A, AH2. and Hz us a standard hydrogen electrode, see cell III described below, E*, electrode potential of the back of a platinum foil in cell III in a solution free of A and AH, us a standard hydrogen electrode, values of E and E* under steady-state conditions of catalytic E,,, &t*, hydrogenation without polarizing current, E eq',equilibrium potential for the electrode reaction H, z 2 H+ + 2 e-, A + E"w 9 equilibrium potential for the electrode reaction AH,. 2 H+ + 2 e-, F, Faraday constant, AG, increase in the Gibbs energy of the reaction A + H, + AH,, I, current applied for polarization (with positive sign for the anodic direction), R, universal gas constant, S, surface area of catalyst, she, standard hydrogen electrode, T, temperature, a,, hydrogen-atom activity in platinum foil, activity in platinum foil under steady-state aH(st) 9 hydrogen-atom conditions of hydrogenation without polarization, au(eq), hydrogen-atom activity in platinum foil in equilibrium with molecular hydrogen H,(aq), a,+, hydrogen-ion activity in the solution,
Hydrogenation of organic compounds in aqueous solutions on noble metals
989
i = I/S, cd applied for polarization, iocjr, exchange cd of reaction j , hAHa, rate of formation of AH, in mol per unit time at surface area S, of compound A, _-ziA, rate of consumption -nHI, rate of hydrogen consumption, &&, hydrogen partial pressure above the solution, st, subscript in order to indicate steady-state conditions without polarizing current (I = 0), Uj, net rate of reaction j in mol per unit area per unit time, aI, charge transfer coefficient in the anodic direction of reaction j, rln = E - E*/, deviation of the electrode potential of the platinum foil from the equilibrium potential of the A-AH, couple. GENERAL
CONSIDERATIONS
In order to differentiate among mechanisms I and IIa or IIb, one may investigate first the rate of the over-all reaction A(aq)
+ Hs(aq) -
AHs(aq)
(16)
as a function of the concentrations of reactants, products and hydrogen ions, which occur as intermediates in the case of mechanism IIa and IIb. Secondly, one may measure the steady-state electrode potential and the steady-state hydrogen-atom activity in the noble metal used as catalyst as a function of the concentrations of reactants, products and hydrogen ions while the reaction A + H, = AH, proceeds. Thirdly, one may investigate the polarization characteristics of the noble metal electrode. Fourthly, one may enforce the equilibrium potential Ees’ or Eeq” on the catalyst as an electrode in an auxiliary circuit with the help of a potentiostat and one may measure (a) the required current I for keeping E = JZ&.,’or Eeqn, respectively, and (b) the rate of consumption of H, or compound A. Fifthly, one may investigate transient phenomena after a sudden change of the potential, or after enforcing suddenly a fmite current upon the noble metal as electrode. In what follows it will be shown that pertinent information may be obtained especially from steady-state potential measurements combined with polarization measurements and, furthermore, measurements of current I and the rate of consumption of the reactant H, or A upon enforcing E = Eeq’, or Eeqn, or E*. POTENTIAL
MEASUREMENTS
The equilibrium potentials Eeq’ and Eecl” corresponding H,(g) *
AH&q) respectively,
to the reactions
2 H+(aq) + 2 e-,
7-t A(aq) + 2 H+(aq) + 2 e-,
(17) w
are obtained by measuring the emf. of the cells
she IIH+(aq), f&&q) she II G+Caq),
If di%rences the dependence
AWaq),
1 Pt,
(1)
A@q) [ Pt.
(11)
between activity and concentration of A and AH, are disregarded, of the single electrode potentials Ees’ and Ee,,” for reactions (17)
C. WAGNER
990
and (18) on the concentrations
of A, AH,, H+ and the partial pressurep3, is given by
Eccl ’
(19)
E aq” = E./O
+
RT z
In
[A]-(au+)s
(20)
[AH,1
The change in the Gibbs energy of the over-all reaction A + Hz = AH, is
AG = --2(~!?,~” - E,,‘)i? The value of AG must be negative when hydrogenation
(21) of A is taking place.
Ees’ -=IEeq”.
Thus
m
The steady-state electrode potential of the catalyst is defined as the potential difference between the metal immersed in the solution involving A, AH,, H, molecules and H+ ions and a standard hydrogen electrode. Despite the occurrence of consecutive anodic and cathodic reactions the net current passing the metal/solution interface is nil if the noble metal is not connected with an outer circuit in accord with general practice during catalytic hydrogenation. The steady-state potential Est in a solution of A, AH,, Hs and hydrogen ions depends on (1) the steady-state activity au(&) of adsorbed hydrogen atoms, (2) the activity of hydrogen ions in the solution, and (3) the net rate 0, of the rate of electrochemical reaction (‘7). The steady-state potential Est is expected to be more noble than the hydrogen equilibrium potential Ees’ for two reasons. First, while the reaction A + Hs = AH, proceeds, the steady-state activity au(&) of adsorbed hydrogen atoms is lower than the activity a=(&) in equilibrium with Hz molecules because of the consumption of adsorbed hydrogen atoms by mechanisms (I), (IIa) and (IIb), which are in parallel. Secondly, if the net rate ZJ~of the electrochemical reaction (7) is positive, the steady-state potential & must be more positive than the potential E,,* , which denotes the equilibrium potential pertaining to the steady-state activity aEfst) of adsorbed hydrogen atoms and the hydrogen-ion concentration in the solution. Thus (23)
Eeq’ < Eat* < EBt < Eeq”.
Previous authorsGB have already measured and evaluated the steady-state potential E,,. At least in some cases, more defkrite conclusions may be drawn upon measuring and evaluating both E,, and E&*. For these measurements, one may use a catalyst, eg platinum in the form of a thin foil, in the double cell she 1
H+(aq) A(aq), AH,(aq) Waq)
Pt(+H)
H+(aq)
she 2.
(III)
One side of the foil is in contact with the solution of A, AH,, H,, and I-I+ in which the reaction A + H, - AH, proceeds. The other side of the foil is in contact with a small volume of solution involving the same hydrogen-ion concentration without the presence of A and AH,. Thus at the latter side the rate of reaction (7) vanishes under steady-state conditions. By virtue of diffusion of hydrogen atoms one obtains the
Hydrogenation
of organic
compoundsin aqueous solutions on noble metals
991
same steady-state hydrogen-atom concentration throughout the foil after the corresponding Ha concentration has been built up in the solution on the right-hand side of cell III. With the help of the double cell III, one may measure ESb as the potential difference between Pt and she 1 and E,,* as the potential difference between Pt and she 2 if I = 0. Instead of standard hydrogen electrodes, one may use suitable reference electrodes in actual experiments in order to facilitate experimentation. The potential B* is given by
where aH is the hydrogen activity in the platinum foi1, which is set equal to unity by defmition for a platinum foil equilibrated with gaseous molecular hydrogen of atmospheric pressure. If equilibrium for reactions (1) and (2) is established corresponding to a hydrogen atom activity +(*a) and the rate of reaction (7) is nil, one has
E eq ’
(25)
From (24) and (25) it follows that for steady-state conditions uHbt)
-
aH(eq)
=
exP
-.
Get*
(26)
Thus a measurement of the difference Emt* - E eq’ yields directly the steady-state hydrogen atom activity in the cataIyst. POLARIZATION
MEASUREMENTS
Further information may be obtained from polarization measurements. For this purpose the noble metal in the form of a foil is connected with an outer electrical circuit including an auxiliary electrode and a reference electrode so that one can pass a well defined net current I across the metal/solution interface, which is given a positive sign for the anodic direction. Under the usual conditions of catalytic hydrogenation, one has I = 0, ie, an equivalence of anodic and cathodic reactions. Upon passing a pre-determined current, one may observe the resulting changes in E and E* with respect to their steady-state values E,$ and Eet*, respectively. Denoting the net rate of reaction j in mol per unit area per unit time by v,, one has ilF = v7 =
07 -
(up + s3 xv,
+
q,)
-
(ur2 + v13 =
07
-
2(Vll
+
~~J,
(27)
since vB = vll and v12= v14. In what follows it is assumed that E”BQ - Eeq'> RT/F. (243) For many reactions of practical interest, (28) holds. Thus the introduction of (28) is not an important restriction of the following considerations. A consequence of (28) is that - AG > RT and, therefore, the over-all reaction is practically irreversible.
C. WAGNER
992
Consequently, if there is a single rate-determining step, it is necessary to consider only the rate of the forward reaction while the rate of the backward reaction can be disregarded. We may envisage the following principal cases: (Ees” (&
-
&)
< (Gl”
-
&I97 < (J%cl” -
R,‘);
(29)
J%sl).
(30)
(1) If E,, is close to Eeql, in accord with (29) and further (Kg” -
E,,) < RT/F,
(31)
the evaluation
of polarization curves is especially straightforward. Suppose that concentration polarization for the molecular species A, A?&, and H+ is negligible. Then, at low pEI values, one may write according to Vetted
where 7” = E - Ees” is the overpotential, jots) and iolll) are the exchange cds of the electron-transfer reactions (9) and (1 I), and a9 and alIare the corresponding chargetransfer coefficients. The negative sign on the right-hand side of (32) occurs since z+,and vu denote the net rate of reactions (9) and (11) in the cathodic direction. If IqF/RTI < 1, one may replace the exponential functions in the numerator of (32) by the first two terms of a series expansion and the exponential functions in the denominator by unity. Then
2~o(B)~om,. VL” - E)F if
t&= t'll = F(&B)
+
E w Ees”.
RT
~OCUJ
(33)
At high pH values one has similarly qg = VI4 =
=
(%I
=-F 2s
+
E)F
~O(9~~O~ll~
&d0(14) +
+
iOCl1,
E w Eea”.
(34)
due to mechanism II at a non-polarized
u14>s
iOf
if
RT
FC&m) + iou4J
Hence, the rate of formation of AH, platinum foil (I = 0; E = &J is %kH,@)
-
2iom,~o(l4) , &”
iom
+
- &IF 1&x,” RT
iO(14)
-
(35)
In order to evaluate (35) one has to determine the expression in brackets with the help of polarization measurements. According to Wagner and Traud,’ the condition EBt = Eeq” requires that in a diagram I vx E the anodic polarization curve is much steeper than the cathodic nolarization curve, ie a(% +
VlZ) =
_2
WJll
+
i+E
%a)
*
(36)
Hydrogenation of organic compounds in aqueous solutions on noble metals
Thus, substituting (33) and (34) in (27) and differentiating one obtains in view of (36)
+
h)io(14) hma,
Substitution
+
h14)
with respect
993
to E,
1.
(37)
of (37) in (35) with i = I/S yields if
Ees” -
RT Es+, < p
,
(38)
Upondividing li an,(II), obtained with the help of (38), by the total rate of formation chemical analysis of the solution in suitable intervals of time, one may calculate the fraction of AH, hydrogenated by consecutive anodic and cathodic reactions instrumental in mechanisms IIa and IIb. If this fraction is close to unity, the electrochemical mechanism prevails. Alternatively, one may polarize the platinum foil to E = Eep” and determine (1) the required current 1 and (2) the rate of consumption of A or formation of AH, for quasi-differential changes in the concentrations of A and AH, so that the shift of E,,’ according to (20) is small in comparison to Eeq” - Est. At E = Eeq” the electrochemical reduction is suppressed. Thus a finite rate of formation of AH, at E = Eecl” mechanism I. is directly a measure of the rate riAn,(I) due to th e non-electrochemical This method applies also if (Ees” - E,,) is higher than RT/F but still (EW” -
E,,) -=z (E,, -
EeJ
Furthermore, if the electrochemical mechanism prevails, the anodic current I required for keeping E = Eeq ” is essentially equivalent to the net rate of the reaction A + H, = AH,atI=O. Thus one has the following criteria in order to find out whether the chemical or the non-electrochemical mechanism prevails. (a) Prevalence of the non-electrochemical mechanism I is indicated if one finds that r,=,eg-i2r:
%
-hH&t)
(3%
and/or (-+rih)E=E,B” = (b) Conversely, one finds that
prevalence
-+A(&)
=
of the electrochemical LEJ~F
w (-&,)se
-fiH&t).
mechanism
(40)
II is indicated
if
(41)
and/or (-7iAh=E,,”
<
(-+H&
(42)
Since for pHz = 1 atm the concentration of molecular hydrogen in an aqueous solution is in general much less than the concentration of the compound A, the over-all rate at I = 0 may be controlled by diffusion of hydrogen to platinum. For example, at a rotating disk the rate may be proportional to the bulk concentration of
c.
994
WAGNER
II, and the square root of the rotational speed@ but independent of the concentrations of A, AH, and H+. Then, according to (27), i/F as a function of E is given by the sum of the constant term v7 and the net cd used for electrochemical reduction of A in a solution containing virtually no hydrogen. Then, regardless of the magnitude of E” - Eet, one may polarize the platinum foil in a solution containing A and AH, $hout hydrogen to the steady-state potential EBbobserved in a solution containing A, AH, and Ha and measure the resulting current. If the observed current I is equivalent to the amount of AII, formed in the solution containing A, AH, and H, without passing current, ie, if I = 2J’(fidH&, = 2I7(---riu&, one may conclude that the electrochemical mechanism IIa or IIb prevails. Conversely, if 1 < 2F(7iliHJst, the non-electrochemical mechanism must prevail. Measurements of E,,* are irrelevant for case 1. (2) If&t - &a’, one does not have the counterpart to the foregoing case E,, M Eecl” because reactions (I) and (2) are instrumental for both the chemical mechanism I and the electrochemical mechanism II. Thus a more detailed discussion is required. At open circuit, the rate of formation of hydrogen atoms due to reaction (2), 2v, = -2tiuJS equals the rate of hydrogen atom consumption due to reactions (4), (5) and (7), -22A&
= 2v, = v, + v6 + v7 = 2v, + v,
if
E = Eat,I = 0.
(43)
If I = 0, the rate of hydrogen consumption equals the rate of formation of AHI. In the case of a polarizing current I applied in the anodic reaction, one has -AuI
= ?i*up -
IJ2E
(44)
In order to suppress anodic ionization of hydrogen, one has to polarize the platinum foil in cell III cathodically so that the front-side and the back-side potential become equal to each other, E = E*. Since E - E* is the driving force for the ionization of hydrogen, v7 vanishes if E = E*. Hence, hydrogen consumption at E = E* is exclusively due to the non-electrochemical mechanism (I), -A,,
= -&*(r)
if
E = E*.
(45)
In view of the consumption of hydrogen atoms by reactions (4) and (5), au at E = E* is lower than au&) and, therefore, according to (24), E* > Eeq’
if
E = E*.
(46)
Thus, applying a gradually increasing cathodic current (-I) to the Pt foil in cell III, one reaches first a state where E = E* > Ees’. Subsequently one reaches a state where E = Eeq’, aH < aH(eq) and E* > Eeq’. This is shown schematically in Figs. 1 and 2 for various conditions. The curves in Fig. 1 refer to conditions where for 1= 0 one has nearly equilibrium for reactions (1) and (2) and the deviation of E from Eeq’ is mainly due to the finite rate constant of reaction (7), ie (&*
-
E,,‘)
< 6% -
whereas the curves in Fig. 2 refer to conditions reaction (7) prevails, ie (Es, -
&*I
< (&*
-
&*I,
where virtually equilibrium J%,J
(47) for (4)
Hydrogenation of organic compounds in
-I
-z
FIG. 1. Plots E and E* UScathodic current (-1) if at I = 0 equilibrium for reactions (1) and (2) is virtually established. (a) The non-ektrochemical mechanism prevails. It>) The electrochemical mechanism prevails, (a)
(b)
-I
-I
FIG. 2. Plots E and E* m cathodic current (-1)
if at I = 0 equilibrium for reaction (7) is virtually established. (a) The non-electrochemical mechanism prevails. (b) The electrochemical mechanism prevails.
Other characteristics of the curves in Figs. 1 and 2 depend on whether the nonelectrochemical or the electrochemical mechanism prevails. (a) If the non-electrochemical mechanism prevails, the hydrogen activity in the catalyst is determined in essence by the interplay of reactions (2), (4) and (5). Thus one has instead of (43) -2riuJS
= 221, w up + v6 = 2v4 > v,.
(4% In the range Ees’ < E < E,,, au and E* are, therefore, agected by cathodic polarization only to a minor extent. Accordingly, the curves E* vs -1in Figs. la and 2a are nearly parallel to the abscissa. If the ionization of hydrogen is suppressed by making E = E* , the rate of hydrogen consumption is decreased, though only slightly since v7 < 2vs at I = 0. Thus ( -~&C==,*
m -fiH*(Bt)*
(50)
C. WAGNER
996
whereupon
in view of (44) (51)
If ES,* is close to E,, ’ in accord with (47, the cathodic current required for enforcing E = Eeg ’ is only slightly higher than that for enforcing E = E*, see Fig. la. Thus equations for -fin, and -1/2F at E = Eeq’ are analogous to (50) and (51). Hence if
(&*
-
J&s’) +Z 6%. -
&*).
(52)
If, however, ESt* is close to Est, no deGnite statement regarding --firr8 and --1/2F at E = Ees’ can be made, since according to Fig. 2a the current required for enforcing E = Ees’ is much greater than that for enforcing E = E*. Adsorbed hydrogen atoms produced by the backward reaction (7) are consumed by reactions (3) and (4). Under these conditions an at E = Eea’ may be considerably greater than an(.t) and accordingly the net rate of reactions (1) and (2) may be lowered in comparison to steady-state conditions. Therefore, -fina at E = Ees’ may be significantly or even much lower than -Ans(stj. (b) If the electrochemical mechanism prevails, the rate of hydrogen consumption at E = E* nearly vanishes, since v, = 0 and -&,(I) is only minor. Thus the curves for E and E* in Figs. la and lb intersect each other slightly above the dotted line E = Ees’. Only a small additional increase of the cathodic current is required in order to enforce E = Ees’. Hence (53) If E - Ees’ ) RTIF, cathodic reduction of A at E = & may be enhanced to a significant extent in comparison to steady-state conditions. Therefore, the current required for keeping E = E* or E = Eeq ’ is supposed to be greater than the equivalent of the current corresponding to (-fint[&& Thus (-~/2JLW (--1/2F)E,Ee(l, 1 ’
-AH&W
(54)
To summarize, if one finds that -tin, at E = E* is close to -&rS(atj and/or -1/2F at E = E* is much less than -riH,{&) , prevalence of the non-electrochemical mechanism is indicated, see (50) and (51). Conversely, if one finds that -fin, at E = B* is much less than --P&,,, and/or --1/2F at E = E* is nearly equal to -i~~,(~~) or greater than -AHl(,ej, p revalence of the electrochemical mechanism is indicated, see (53) and (54). Conclusions, however, are less definite if one tries to evaluate data for E = ,?&’ rather than for E = &*. If one finds that -finI at E = Eeq’ is close to -?&(#tj and/or -1/2F at E = Ees’ is much less than -P!Z~*(~) p revalence of the non-electrochemical mechanism is indicated. On the other hand, if one finds that --An, at E = Eeq’ is much less than -&r,(st), no definite conclusions can be drawn since this result can be expected not only in the case of prevalence of the electrochemical mechanism but also in the case of prevalence of the non-electrochemical mechanism, see the discussion following (52).
Hydrogenationof organic compounds in aqueous solutionson noble metals
997
Wagner and Traud’ investigated the mechanism of the catalytic reduction of by H, on a Pt foil as catalyst by measuring -zi,+ at C,H,NO,, S,O?- and H&O, an enforced potential E = Eeq'.Since the rate of hydrogen consumption at E = Eeq' was practically nil, it was concluded in accord with (53) that the electrochemical mechanism prevails. This conclusion, however, is not definite since negligible hydrogen consumption at E = Ees'may also be found if the non-electrochemical mechanism prevails. The conclusion drawn by Wagner and Traud’ would be justified if it can be shown that polarization of the reaction & H, = H+(aq) + e- is mainly due to reaction (7). This supposition is in accord with an evaluation of polarization curves in solutions free of A and AH, over a fairly wide potential range. 6~10 It is likely but not certain To remove that this is also true in solutions containing C,H,NO,, S,082- or H,AsO,. possible objections to this hypothesis and to obtain direct information, it is desirable to measure -tiiH2 at E = E* in solutions of C6H,N02, S,0s2-, or H3As04 and to apply the criteria in (50) and (51). CONCLUDING
REMARKS
The objective of the measurements considered above is to obtain criteria in order to distinguish between prevalence of the non-electrochemical mechanism and prevalence of the electrochemical mechanism. If such a decision has been reached, it is desirable to ascertain further details of the prevailing mechanism with the help of other measurements. Such a discussion, however, is beyond the scope of this paper. In contradistinction to traditional reasoning in reaction kinetics, the dependence of the reaction rate on the concentrations of reactants and products is not considered in this paper. Definite conclusions are possible only if additional information regarding the degree of coverage of the Pt surface by the various species is available. It is important to notice that the evaluation of measurements suggested in this paper is independent of special assumptions on the composition of the adsorption layer. Experimental investigations in accord with the foregoing considerations have been conducted by Takehara and are reported in a following paper.ll Acknowledgements-The
criticalcomments.
author wishesto thank ProfessorZ. Takehara for helpful discussionsand REFERENCES
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Acta. 1970. Vol. 15, pp. 987 to 997.
Pcrktunon Press. Printed in Northern Ireland
CONSIDERATIONS ON THE MECHANISM OF THE HYDROGENATION OF ORGANIC COMPOUNDS IN AQUEOUS SOLUTIONS ON NOBLE METAL CATALYSTS* C. WAGNER Max-Planck-lnstitut fllr physikahsche Chemie, GGttingen, Bundesrepublik, Deutschland Abstract-Hydrogenation of an organic compound A on Pt as catalyst in aqueous solutions may occur (1) by a sequence of non-electiochemical reactions involving adsorbed hydrogen atoms and radicals as intermediates and (2) by anodic ionization of hydrogen and consecutive cathodic reduction of compound A. This paper is concerned with experimental methods that are especially promising for deciding whether a non-electrochemical or an electrochemical mechanismprevails for a particular compound. R&um&L’hydrog&ration d’un compose organique A en solutions aqueuses, avec le Pt comme catalyseur, peut se manifester comme une succession de reactions non electrochimiqu~, impliquant comme intermecliaires,Padsorption d’atomes d’hydro@ne et de radicaux, l’ionisation anodique de l’hydrogene et la reduction cathodique consecutive du compose A. Le present memoire disc&e ies methodes experimentales les plus inttkessantes,en vue de decider si urt type de m&anisme, non6lectrochimique ou &ctrochimique, prevaut pour un cornpod particulier. Zussammenfassrmg-Die Hydrierung einer organischen Verbindung A an Pt als Katalysator karm entweder durch eine Folge nichtelektrochemischer Reaktionen mit adsorbierten Wasserstoffatomen .und Radikalen als Zwischenprodukt oder durch anodische Ionisierung von WasserstofI und kathodische Reduktion der Verbmdung A erfolgen. In dieser Arbeit werden experimentelle Methoden angegeben, die besonders gee&net scheinen, um zu entscheiden, ob em nichtelektrochemischer oder ein elektrochemischer Mechanismus im Einzelfall vorliegt. STATEMENT
OF THB
PROBLEM
THE HYDROGENATION of an organic compound A by H, to AH, in an aqueous solution
with the help of a noble metal, eg platinum, as a catalyst, may occur via various mechanisms.lve First, hydrogen molecules may dissociate on the surface of the noble metal and subsequently hydrogen atoms may react with compound A in consecutive steps in the same way as the catalytic hydrogenation of a gaseous organic compound, eg ethylene at a metal/gas interface proceeds with adsorbed hydrogen atoms and adsorbed AH radicals as intermediates (mechanism I) Hs(aq) H,(ad) A(aq) -
Ha(ad), 2 II(
(2)
A(ad),
(3)
A(ad) + H(ad) - AH(ad), AH(ad) + H(ad) - AH,(ad),
(1)
(4) (5)
AH&d) - AH&q). (6) Secondly, electrochemical mechanisms involving consecutive anodic and cathodic reactions with zero net current at the metal/solution interface may be considered. Investigations on the cathodic reduction of quinone to hydroquinone by Vetterf have shown that different electron-transfer reactions prevail at low pH values (pH < 5) and high pH values (pH > 6). In view of these results one may consider the following sequences of reactions with prevalence of reactions on the left-hand side (mechanism IIa) at low pII values and prevalence of reactions on the right-hand side * Manuscript received 10 April 1969. 987
C. WAGNER
988
{mechanism
IIb) at high pH values, H,(aq) -+ H,(ad), H,(ad) - 2 H(ad), A(aq) - A(ad), 2(H(ad) -+ H+(aq)
A(ad) + H+(aq) ---f AH+(ad), AH+(acl) + e- AH(ad)
+ H+(aq) -
AH$(ad)
+ e- -
(1) (2) (3) + e-1;
AH(ad),
A-(ad)
AHz+(ad),
(9) (10)
AH,(ad) ,
(11)
AH-(ad)
AH,(ad)
(7)
A(ad) + e- -
(8)
+ H+(aq) -
A-(ad),
(12)
AH(ad),
+ e- -
AH-(ad),
(13) (14)
+ H+(aq) -
AH,(ad),
(15)
AH(ad) -+ AH,(aq).
(6) One may hypothesize that mechanism I prevails in the case of hydrogenation of unsaturated hydrocarbons and conversely electrochemical mechanism (IIa) or (IIb) prevails in systems where the equilibrium potential of the redox couple A-AH, at a noble metal is readily established, ie the exchange cd of the couple A-AH, is relatively high, as in the case of the quinone-hydroquinone couple. It is the objective of this paper to outline experimental investigations which are promising in order to decide which mechanism is instrumental for hydrogenation of a particular inorganic or organic compound A. SYMBOLS
We use the following: E, electrode potential of a platinum .foil in a solution containing A, AH2. and Hz us a standard hydrogen electrode, see cell III described below, E*, electrode potential of the back of a platinum foil in cell III in a solution free of A and AH, us a standard hydrogen electrode, values of E and E* under steady-state conditions of catalytic E,,, &t*, hydrogenation without polarizing current, E eq',equilibrium potential for the electrode reaction H, z 2 H+ + 2 e-, A + E"w 9 equilibrium potential for the electrode reaction AH,. 2 H+ + 2 e-, F, Faraday constant, AG, increase in the Gibbs energy of the reaction A + H, + AH,, I, current applied for polarization (with positive sign for the anodic direction), R, universal gas constant, S, surface area of catalyst, she, standard hydrogen electrode, T, temperature, a,, hydrogen-atom activity in platinum foil, activity in platinum foil under steady-state aH(st) 9 hydrogen-atom conditions of hydrogenation without polarization, au(eq), hydrogen-atom activity in platinum foil in equilibrium with molecular hydrogen H,(aq), a,+, hydrogen-ion activity in the solution,
Hydrogenation of organic compounds in aqueous solutions on noble metals
989
i = I/S, cd applied for polarization, iocjr, exchange cd of reaction j , hAHa, rate of formation of AH, in mol per unit time at surface area S, of compound A, _-ziA, rate of consumption -nHI, rate of hydrogen consumption, &&, hydrogen partial pressure above the solution, st, subscript in order to indicate steady-state conditions without polarizing current (I = 0), Uj, net rate of reaction j in mol per unit area per unit time, aI, charge transfer coefficient in the anodic direction of reaction j, rln = E - E*/, deviation of the electrode potential of the platinum foil from the equilibrium potential of the A-AH, couple. GENERAL
CONSIDERATIONS
In order to differentiate among mechanisms I and IIa or IIb, one may investigate first the rate of the over-all reaction A(aq)
+ Hs(aq) -
AHs(aq)
(16)
as a function of the concentrations of reactants, products and hydrogen ions, which occur as intermediates in the case of mechanism IIa and IIb. Secondly, one may measure the steady-state electrode potential and the steady-state hydrogen-atom activity in the noble metal used as catalyst as a function of the concentrations of reactants, products and hydrogen ions while the reaction A + H, = AH, proceeds. Thirdly, one may investigate the polarization characteristics of the noble metal electrode. Fourthly, one may enforce the equilibrium potential Ees’ or Eeq” on the catalyst as an electrode in an auxiliary circuit with the help of a potentiostat and one may measure (a) the required current I for keeping E = JZ&.,’or Eeqn, respectively, and (b) the rate of consumption of H, or compound A. Fifthly, one may investigate transient phenomena after a sudden change of the potential, or after enforcing suddenly a fmite current upon the noble metal as electrode. In what follows it will be shown that pertinent information may be obtained especially from steady-state potential measurements combined with polarization measurements and, furthermore, measurements of current I and the rate of consumption of the reactant H, or A upon enforcing E = Eeq’, or Eeqn, or E*. POTENTIAL
MEASUREMENTS
The equilibrium potentials Eeq’ and Eecl” corresponding H,(g) *
AH&q) respectively,
to the reactions
2 H+(aq) + 2 e-,
7-t A(aq) + 2 H+(aq) + 2 e-,
(17) w
are obtained by measuring the emf. of the cells
she IIH+(aq), f&&q) she II G+Caq),
If di%rences the dependence
AWaq),
1 Pt,
(1)
A@q) [ Pt.
(11)
between activity and concentration of A and AH, are disregarded, of the single electrode potentials Ees’ and Ee,,” for reactions (17)
C. WAGNER
990
and (18) on the concentrations
of A, AH,, H+ and the partial pressurep3, is given by
Eccl ’
(19)
E aq” = E./O
+
RT z
In
[A]-(au+)s
(20)
[AH,1
The change in the Gibbs energy of the over-all reaction A + Hz = AH, is
AG = --2(~!?,~” - E,,‘)i? The value of AG must be negative when hydrogenation
(21) of A is taking place.
Ees’ -=IEeq”.
Thus
m
The steady-state electrode potential of the catalyst is defined as the potential difference between the metal immersed in the solution involving A, AH,, H, molecules and H+ ions and a standard hydrogen electrode. Despite the occurrence of consecutive anodic and cathodic reactions the net current passing the metal/solution interface is nil if the noble metal is not connected with an outer circuit in accord with general practice during catalytic hydrogenation. The steady-state potential Est in a solution of A, AH,, Hs and hydrogen ions depends on (1) the steady-state activity au(&) of adsorbed hydrogen atoms, (2) the activity of hydrogen ions in the solution, and (3) the net rate 0, of the rate of electrochemical reaction (‘7). The steady-state potential Est is expected to be more noble than the hydrogen equilibrium potential Ees’ for two reasons. First, while the reaction A + Hs = AH, proceeds, the steady-state activity au(&) of adsorbed hydrogen atoms is lower than the activity a=(&) in equilibrium with Hz molecules because of the consumption of adsorbed hydrogen atoms by mechanisms (I), (IIa) and (IIb), which are in parallel. Secondly, if the net rate ZJ~of the electrochemical reaction (7) is positive, the steady-state potential & must be more positive than the potential E,,* , which denotes the equilibrium potential pertaining to the steady-state activity aEfst) of adsorbed hydrogen atoms and the hydrogen-ion concentration in the solution. Thus (23)
Eeq’ < Eat* < EBt < Eeq”.
Previous authorsGB have already measured and evaluated the steady-state potential E,,. At least in some cases, more defkrite conclusions may be drawn upon measuring and evaluating both E,, and E&*. For these measurements, one may use a catalyst, eg platinum in the form of a thin foil, in the double cell she 1
H+(aq) A(aq), AH,(aq) Waq)
Pt(+H)
H+(aq)
she 2.
(III)
One side of the foil is in contact with the solution of A, AH,, H,, and I-I+ in which the reaction A + H, - AH, proceeds. The other side of the foil is in contact with a small volume of solution involving the same hydrogen-ion concentration without the presence of A and AH,. Thus at the latter side the rate of reaction (7) vanishes under steady-state conditions. By virtue of diffusion of hydrogen atoms one obtains the
Hydrogenation
of organic
compoundsin aqueous solutions on noble metals
991
same steady-state hydrogen-atom concentration throughout the foil after the corresponding Ha concentration has been built up in the solution on the right-hand side of cell III. With the help of the double cell III, one may measure ESb as the potential difference between Pt and she 1 and E,,* as the potential difference between Pt and she 2 if I = 0. Instead of standard hydrogen electrodes, one may use suitable reference electrodes in actual experiments in order to facilitate experimentation. The potential B* is given by
where aH is the hydrogen activity in the platinum foi1, which is set equal to unity by defmition for a platinum foil equilibrated with gaseous molecular hydrogen of atmospheric pressure. If equilibrium for reactions (1) and (2) is established corresponding to a hydrogen atom activity +(*a) and the rate of reaction (7) is nil, one has
E eq ’
(25)
From (24) and (25) it follows that for steady-state conditions uHbt)
-
aH(eq)
=
exP
-.
Get*
(26)
Thus a measurement of the difference Emt* - E eq’ yields directly the steady-state hydrogen atom activity in the cataIyst. POLARIZATION
MEASUREMENTS
Further information may be obtained from polarization measurements. For this purpose the noble metal in the form of a foil is connected with an outer electrical circuit including an auxiliary electrode and a reference electrode so that one can pass a well defined net current I across the metal/solution interface, which is given a positive sign for the anodic direction. Under the usual conditions of catalytic hydrogenation, one has I = 0, ie, an equivalence of anodic and cathodic reactions. Upon passing a pre-determined current, one may observe the resulting changes in E and E* with respect to their steady-state values E,$ and Eet*, respectively. Denoting the net rate of reaction j in mol per unit area per unit time by v,, one has ilF = v7 =
07 -
(up + s3 xv,
+
q,)
-
(ur2 + v13 =
07
-
2(Vll
+
~~J,
(27)
since vB = vll and v12= v14. In what follows it is assumed that E”BQ - Eeq'> RT/F. (243) For many reactions of practical interest, (28) holds. Thus the introduction of (28) is not an important restriction of the following considerations. A consequence of (28) is that - AG > RT and, therefore, the over-all reaction is practically irreversible.
C. WAGNER
992
Consequently, if there is a single rate-determining step, it is necessary to consider only the rate of the forward reaction while the rate of the backward reaction can be disregarded. We may envisage the following principal cases: (Ees” (&
-
&)
< (Gl”
-
&I97 < (J%cl” -
R,‘);
(29)
J%sl).
(30)
(1) If E,, is close to Eeql, in accord with (29) and further (Kg” -
E,,) < RT/F,
(31)
the evaluation
of polarization curves is especially straightforward. Suppose that concentration polarization for the molecular species A, A?&, and H+ is negligible. Then, at low pEI values, one may write according to Vetted
where 7” = E - Ees” is the overpotential, jots) and iolll) are the exchange cds of the electron-transfer reactions (9) and (1 I), and a9 and alIare the corresponding chargetransfer coefficients. The negative sign on the right-hand side of (32) occurs since z+,and vu denote the net rate of reactions (9) and (11) in the cathodic direction. If IqF/RTI < 1, one may replace the exponential functions in the numerator of (32) by the first two terms of a series expansion and the exponential functions in the denominator by unity. Then
2~o(B)~om,. VL” - E)F if
t&= t'll = F(&B)
+
E w Ees”.
RT
~OCUJ
(33)
At high pH values one has similarly qg = VI4 =
=
(%I
=-F 2s
+
E)F
~O(9~~O~ll~
&d0(14) +
+
iOCl1,
E w Eea”.
(34)
due to mechanism II at a non-polarized
u14>s
iOf
if
RT
FC&m) + iou4J
Hence, the rate of formation of AH, platinum foil (I = 0; E = &J is %kH,@)
-
2iom,~o(l4) , &”
iom
+
- &IF 1&x,” RT
iO(14)
-
(35)
In order to evaluate (35) one has to determine the expression in brackets with the help of polarization measurements. According to Wagner and Traud,’ the condition EBt = Eeq” requires that in a diagram I vx E the anodic polarization curve is much steeper than the cathodic nolarization curve, ie a(% +
VlZ) =
_2
WJll
+
i+E
%a)
*
(36)
Hydrogenation of organic compounds in aqueous solutions on noble metals
Thus, substituting (33) and (34) in (27) and differentiating one obtains in view of (36)
+
h)io(14) hma,
Substitution
+
h14)
with respect
993
to E,
1.
(37)
of (37) in (35) with i = I/S yields if
Ees” -
RT Es+, < p
,
(38)
Upondividing li an,(II), obtained with the help of (38), by the total rate of formation chemical analysis of the solution in suitable intervals of time, one may calculate the fraction of AH, hydrogenated by consecutive anodic and cathodic reactions instrumental in mechanisms IIa and IIb. If this fraction is close to unity, the electrochemical mechanism prevails. Alternatively, one may polarize the platinum foil to E = Eep” and determine (1) the required current 1 and (2) the rate of consumption of A or formation of AH, for quasi-differential changes in the concentrations of A and AH, so that the shift of E,,’ according to (20) is small in comparison to Eeq” - Est. At E = Eeq” the electrochemical reduction is suppressed. Thus a finite rate of formation of AH, at E = Eecl” mechanism I. is directly a measure of the rate riAn,(I) due to th e non-electrochemical This method applies also if (Ees” - E,,) is higher than RT/F but still (EW” -
E,,) -=z (E,, -
EeJ
Furthermore, if the electrochemical mechanism prevails, the anodic current I required for keeping E = Eeq ” is essentially equivalent to the net rate of the reaction A + H, = AH,atI=O. Thus one has the following criteria in order to find out whether the chemical or the non-electrochemical mechanism prevails. (a) Prevalence of the non-electrochemical mechanism I is indicated if one finds that r,=,eg-i2r:
%
-hH&t)
(3%
and/or (-+rih)E=E,B” = (b) Conversely, one finds that
prevalence
-+A(&)
=
of the electrochemical LEJ~F
w (-&,)se
-fiH&t).
mechanism
(40)
II is indicated
if
(41)
and/or (-7iAh=E,,”
<
(-+H&
(42)
Since for pHz = 1 atm the concentration of molecular hydrogen in an aqueous solution is in general much less than the concentration of the compound A, the over-all rate at I = 0 may be controlled by diffusion of hydrogen to platinum. For example, at a rotating disk the rate may be proportional to the bulk concentration of
c.
994
WAGNER
II, and the square root of the rotational speed@ but independent of the concentrations of A, AH, and H+. Then, according to (27), i/F as a function of E is given by the sum of the constant term v7 and the net cd used for electrochemical reduction of A in a solution containing virtually no hydrogen. Then, regardless of the magnitude of E” - Eet, one may polarize the platinum foil in a solution containing A and AH, $hout hydrogen to the steady-state potential EBbobserved in a solution containing A, AH, and Ha and measure the resulting current. If the observed current I is equivalent to the amount of AII, formed in the solution containing A, AH, and H, without passing current, ie, if I = 2J’(fidH&, = 2I7(---riu&, one may conclude that the electrochemical mechanism IIa or IIb prevails. Conversely, if 1 < 2F(7iliHJst, the non-electrochemical mechanism must prevail. Measurements of E,,* are irrelevant for case 1. (2) If&t - &a’, one does not have the counterpart to the foregoing case E,, M Eecl” because reactions (I) and (2) are instrumental for both the chemical mechanism I and the electrochemical mechanism II. Thus a more detailed discussion is required. At open circuit, the rate of formation of hydrogen atoms due to reaction (2), 2v, = -2tiuJS equals the rate of hydrogen atom consumption due to reactions (4), (5) and (7), -22A&
= 2v, = v, + v6 + v7 = 2v, + v,
if
E = Eat,I = 0.
(43)
If I = 0, the rate of hydrogen consumption equals the rate of formation of AHI. In the case of a polarizing current I applied in the anodic reaction, one has -AuI
= ?i*up -
IJ2E
(44)
In order to suppress anodic ionization of hydrogen, one has to polarize the platinum foil in cell III cathodically so that the front-side and the back-side potential become equal to each other, E = E*. Since E - E* is the driving force for the ionization of hydrogen, v7 vanishes if E = E*. Hence, hydrogen consumption at E = E* is exclusively due to the non-electrochemical mechanism (I), -A,,
= -&*(r)
if
E = E*.
(45)
In view of the consumption of hydrogen atoms by reactions (4) and (5), au at E = E* is lower than au&) and, therefore, according to (24), E* > Eeq’
if
E = E*.
(46)
Thus, applying a gradually increasing cathodic current (-I) to the Pt foil in cell III, one reaches first a state where E = E* > Ees’. Subsequently one reaches a state where E = Eeq’, aH < aH(eq) and E* > Eeq’. This is shown schematically in Figs. 1 and 2 for various conditions. The curves in Fig. 1 refer to conditions where for 1= 0 one has nearly equilibrium for reactions (1) and (2) and the deviation of E from Eeq’ is mainly due to the finite rate constant of reaction (7), ie (&*
-
E,,‘)
< 6% -
whereas the curves in Fig. 2 refer to conditions reaction (7) prevails, ie (Es, -
&*I
< (&*
-
&*I,
where virtually equilibrium J%,J
(47) for (4)
Hydrogenation of organic compounds in
-I
-z
FIG. 1. Plots E and E* UScathodic current (-1) if at I = 0 equilibrium for reactions (1) and (2) is virtually established. (a) The non-ektrochemical mechanism prevails. It>) The electrochemical mechanism prevails, (a)
(b)
-I
-I
FIG. 2. Plots E and E* m cathodic current (-1)
if at I = 0 equilibrium for reaction (7) is virtually established. (a) The non-electrochemical mechanism prevails. (b) The electrochemical mechanism prevails.
Other characteristics of the curves in Figs. 1 and 2 depend on whether the nonelectrochemical or the electrochemical mechanism prevails. (a) If the non-electrochemical mechanism prevails, the hydrogen activity in the catalyst is determined in essence by the interplay of reactions (2), (4) and (5). Thus one has instead of (43) -2riuJS
= 221, w up + v6 = 2v4 > v,.
(4% In the range Ees’ < E < E,,, au and E* are, therefore, agected by cathodic polarization only to a minor extent. Accordingly, the curves E* vs -1in Figs. la and 2a are nearly parallel to the abscissa. If the ionization of hydrogen is suppressed by making E = E* , the rate of hydrogen consumption is decreased, though only slightly since v7 < 2vs at I = 0. Thus ( -~&C==,*
m -fiH*(Bt)*
(50)
C. WAGNER
996
whereupon
in view of (44) (51)
If ES,* is close to E,, ’ in accord with (47, the cathodic current required for enforcing E = Eeg ’ is only slightly higher than that for enforcing E = E*, see Fig. la. Thus equations for -fin, and -1/2F at E = Eeq’ are analogous to (50) and (51). Hence if
(&*
-
J&s’) +Z 6%. -
&*).
(52)
If, however, ESt* is close to Est, no deGnite statement regarding --firr8 and --1/2F at E = Ees’ can be made, since according to Fig. 2a the current required for enforcing E = Ees’ is much greater than that for enforcing E = E*. Adsorbed hydrogen atoms produced by the backward reaction (7) are consumed by reactions (3) and (4). Under these conditions an at E = Eea’ may be considerably greater than an(.t) and accordingly the net rate of reactions (1) and (2) may be lowered in comparison to steady-state conditions. Therefore, -fina at E = Ees’ may be significantly or even much lower than -Ans(stj. (b) If the electrochemical mechanism prevails, the rate of hydrogen consumption at E = E* nearly vanishes, since v, = 0 and -&,(I) is only minor. Thus the curves for E and E* in Figs. la and lb intersect each other slightly above the dotted line E = Ees’. Only a small additional increase of the cathodic current is required in order to enforce E = Ees’. Hence (53) If E - Ees’ ) RTIF, cathodic reduction of A at E = & may be enhanced to a significant extent in comparison to steady-state conditions. Therefore, the current required for keeping E = E* or E = Eeq ’ is supposed to be greater than the equivalent of the current corresponding to (-fint[&& Thus (-~/2JLW (--1/2F)E,Ee(l, 1 ’
-AH&W
(54)
To summarize, if one finds that -tin, at E = E* is close to -&rS(atj and/or -1/2F at E = E* is much less than -riH,{&) , prevalence of the non-electrochemical mechanism is indicated, see (50) and (51). Conversely, if one finds that -fin, at E = B* is much less than --P&,,, and/or --1/2F at E = E* is nearly equal to -i~~,(~~) or greater than -AHl(,ej, p revalence of the electrochemical mechanism is indicated, see (53) and (54). Conclusions, however, are less definite if one tries to evaluate data for E = ,?&’ rather than for E = &*. If one finds that -finI at E = Eeq’ is close to -?&(#tj and/or -1/2F at E = Ees’ is much less than -P!Z~*(~) p revalence of the non-electrochemical mechanism is indicated. On the other hand, if one finds that --An, at E = Eeq’ is much less than -&r,(st), no definite conclusions can be drawn since this result can be expected not only in the case of prevalence of the electrochemical mechanism but also in the case of prevalence of the non-electrochemical mechanism, see the discussion following (52).
Hydrogenationof organic compounds in aqueous solutionson noble metals
997
Wagner and Traud’ investigated the mechanism of the catalytic reduction of by H, on a Pt foil as catalyst by measuring -zi,+ at C,H,NO,, S,O?- and H&O, an enforced potential E = Eeq'.Since the rate of hydrogen consumption at E = Eeq' was practically nil, it was concluded in accord with (53) that the electrochemical mechanism prevails. This conclusion, however, is not definite since negligible hydrogen consumption at E = Ees'may also be found if the non-electrochemical mechanism prevails. The conclusion drawn by Wagner and Traud’ would be justified if it can be shown that polarization of the reaction & H, = H+(aq) + e- is mainly due to reaction (7). This supposition is in accord with an evaluation of polarization curves in solutions free of A and AH, over a fairly wide potential range. 6~10 It is likely but not certain To remove that this is also true in solutions containing C,H,NO,, S,082- or H,AsO,. possible objections to this hypothesis and to obtain direct information, it is desirable to measure -tiiH2 at E = E* in solutions of C6H,N02, S,0s2-, or H3As04 and to apply the criteria in (50) and (51). CONCLUDING
REMARKS
The objective of the measurements considered above is to obtain criteria in order to distinguish between prevalence of the non-electrochemical mechanism and prevalence of the electrochemical mechanism. If such a decision has been reached, it is desirable to ascertain further details of the prevailing mechanism with the help of other measurements. Such a discussion, however, is beyond the scope of this paper. In contradistinction to traditional reasoning in reaction kinetics, the dependence of the reaction rate on the concentrations of reactants and products is not considered in this paper. Definite conclusions are possible only if additional information regarding the degree of coverage of the Pt surface by the various species is available. It is important to notice that the evaluation of measurements suggested in this paper is independent of special assumptions on the composition of the adsorption layer. Experimental investigations in accord with the foregoing considerations have been conducted by Takehara and are reported in a following paper.ll Acknowledgements-The
criticalcomments.
author wishesto thank ProfessorZ. Takehara for helpful discussionsand REFERENCES
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