Electrocatalytic hydrogenation processes at controlled
potential. 1. The electrocatalytic reduction of nitric acid
Citation for published version (APA):
Plas, van der, J. F., & Barendrecht, E. (1980). Electrocatalytic hydrogenation processes at controlled potential.
1. The electrocatalytic reduction of nitric acid. Electrochimica Acta, 25(11), 1463-1469.
https://doi.org/10.1016/0013-4686(80)87162-3
DOI:
10.1016/0013-4686(80)87162-3
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Published: 01/01/1980
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Elecrrochmica Acfo, Vol. 25, pp. 1463-1469. Pergaman Press Ltd. 1980. Printed ie Great Britain
ELECTROCATALYTIC
HYDROGENATION
PROCESSES
CONTROLLED
POTENTIAL
- 1. THE
ELECTROCATALYTIC
REDUCTION
OF NITRIC ACID
J. F. VAN DER PLAS and E. BARENDRE~HT
Laboratory for Electrochemistry, University of Technology, P.O. Box 513, Eindhoven, The Netherlands (Received 30 April 1979)
Abslract-Methods are described to control the catalyst potential during the electrocatalytic hydrogenation of nitric acid in a slurry electrode cell. The influence of traces of GeO, on the hydrogenation reaction is studied with a platinum electrode. Deposition of Ge on the platinum surface limits the amount of adsorbed hydrogen and enhances the nitric acid reduction. This changes the mixed potential of the hydrogen-nitric
acid svstem on the olatinum (catalvst) surface. A similar potential change occurs, when the partial hydrogen press&e in the sy&n is al&red. _
1. INTRODUCTION
The catalytic hydrogenation of nitric acid to hy- droxylamine is an important industrial process[l-31. The hydroxylamine formed is a valuable intermediate in the manufacturing of some bulk chemicals as caprolactam. The hydrogenation rate can be improved by adding small amounts (about 10m5 moles/g cata- lyst) of a “promoting” metal, mostly as its salt or oxide[3]. To prevent the formation of ammonia, the reduction must be performed in an acid buffer eg a phosphoric acid system, 0 c pH < 3, (a bisulphate buffer can also be used[3]). In a too acidic solution, however, the catalyst (normally a carbon supported metal of the platinum group) can dissolve.
Electrochemically, the reduction of nitric acid re- quires a relatively high overpotential. However, in the presence of nitrous acid the reduction proceeds very smoothly (as was reported already by Abel and Schmid[4-6] in 1928), due to an autocatalyticreaction of nitric acid-nitrate with nitrogen oxide, NO, formed by the electrolytic reduction of nitrous acid. Further reduction of nitrous acid to hydroxylamine at a platinum electrode goes smoothly.
A special feature of the catalytic process, therefore, seems to be the accelerating effect of the catalyst on the reduction rate of nitric acid to nitrous acid. This reaction step is even more enhanced by using a catalyst promotor. It is probable that electrochemical reaction steps are part of the further reduction of nitrous acid. In that case, the potential of the catalyst determines the extent of possible electrochemical reactions and so selectivity controls product formation. Small amounts of catalyst promotor change the surface of the catalyst and so influence its behaviour as an electrode. This effect is known as the result of (underpotential) deposition of submonolayers of metal on an electrode surface[7-lo], thus inhibiting an electrode poisoning reaction. So, simultaneous reactions may occur on the catalyst surface., influencing the catalyst potential. In 1974, Kinza[ 111 investigated the influence of the catalyst potential on the product distribution for the electrocatalytic reduction of nitric acid. The catalyst potential changed as a function of the amount of nitric
acid in the solution ; however, the product distribution did not change. We investigated methods[12,13] to change the catalyst potential and its effect on product distribution.
Because of both the catalytic and the electrochemi- cal nature of the hydrogenation of nitric acid, we chose this reaction as a model to study both the kinetics (mechanism and rate) and the technological impli- cations (mass and charge transfer). The extensive effect of a catalyst promotor (one of the most effective is germanium) forced us to study its influence on the hydrogen adsorption and desorption at a platinum surface and on the electrocatalytic properties of the formed platinum-germanium surface for the reduction of nitric acid.
2. EXPERIMENTAL
The catalytic hydrogenation reactions of nitric acid is studied in a slurry electrode reactor (Fig. l), equipped with a platinum gauze electrode (area 46 cm’) as a working electrode.
I-E curves were recorded using a Wenking poten-
tiostat, type 68 FR 0.5. The potential of the catalyst particles was measured with a special designed measuring probe (Fig. 2), a small platinum wire combined with a reference electrode (saturated mer- cury sulphate).
Experiments were carried out in sulphuric acid-potassium sulphate solutions of different com- position, with nitric acid or sodium nitrate upto 0.25 M. The catalyst was carbon supported platinum (Merck nr. 807339).
The hydrogen (sometimes diluted by argon) was made oxygen-free with a copper catalyst (BASF R3-11).
Promotor experiments were carried out in a glass cetl, thermostatted at 25”C, with a cathode compart- ment (150 ml), separated from the anode compartment by means of a glass frit. Contact with the reference electrode (a Hg; Hg,SO.,, sat. K,SO, ~ half cell, + 0.70 Y us she) was made with a horizontally placed Luggin capillary. The disc electrodes used (area
1464 J. F. VANDER PLAS AND E. BARENDRECHT
Glass fit
Fig. 1. Slurry electrode reactor with feeder electrode; reactor
contents: 250ml; gauze width: 1 mm.
0.3cm’) were mounted in a stainless steel electrode holder, shaped to fit exactly in a NS 29132 standard ground joint. -
Z-E and I-cu relations were obtained with a Tacussel bipotentiostat, type BI-PAD, in combination with a Wenking linear voltage scangenerator, type VSG72. The curves were recorded with a Hewlett Packard X(t)-Y-recorder, type 7046, and the electrode poten- tials were measured with Philips dc-microvoltmeters, PM 2435. The rotation speed of the electrode could be
I i
Reference electrode
cmpmmenr c I
Glass f rity ‘- Measuring probe
Fig. 2. Combined electrode system for the measurement of the catalyst potential; area probe: 0.4cm*.
varied between 100 and 50OOrpm and was kept constant at a desired value by means of a tachogen- erator, type Motomatic. The potential was varied between the potentials where hydrogen, respectively oxygen, are formed at the electrode.
These experiments were performed in a 1.25 M sulphuric acid solution, saturated with nitrogen or hydrogen. A germanium dioxide solution, 5.10m3 M, was added to the solution upto a concentration of 3.10-&M. Nitric acid was added in a concentration of 25. 10e3 M. All chemicals were analytical grade.
3. RESULTS
3.1. Slurry electrode experiments
To study charge transfer, first the current-potential relationship for the oxidation of hydrogen in the sulphuric acid solution with and without the catalyst suspension is determined (Fig. 3). As could be expec- ted, in the presence of the catalyst, the current is higher than in its absence. So, charge is transferred between the catalyst and the feeder electrode. An analogous effect is shown for different rotation speeds of the stirrer (Fig. 4). In the absence of the catalyst, a limiting current is obtained at higher rotation speeds. How- ever, in the presence of the catalyst no limiting current was found in the rotation speed range mea- sured, due to the increasing collision frequency of the catalyst particles with the work electrode. Although it is evident that charge transfer takes place between the working (feeder) electrode and the catalyst particles, no change in the potential of the catalyst particles was measured. So, the use of a feeder electrode to influence the potential of the catalyst particles was not effective. Next, we therefore studied the possibility of in- fluencing the catalyst potential by changing the re- action conditions (the medium), without the working electrode. First, it was found that the potential of the catalyst depends on the rotation speed of the stirrer (Fig. 5). This is due to the inefficiency of dissoluting hydrogen, being not optimal at low rotation speeds. At lower hydrogen pressures a catalyst potential change of several hundreds of millivolts was measured if the hydrogen pressure was reduced by a factor of ten (Fig. 6).
Still larger potential variations were found, if small amounts of germanium dioxide (up to 5mM) were added to the solution. It is known, that germanium accelerates the catalytic hydrogenation of nitric acid for the type ofcatalysts applied[3,14]. These potential changes were also influenced by the changes in hy- drogen pressure (Fig. 6). However, the changes in catalyst potential as a function of the hydrogen pressure were not of permanent character. In most cases, the catalyst potential returned slowly to the value of the reversible hydrogen potential in the solution after about 18 hours. It was therefore not possible to study quantitatively the influence of these potential changes on the reactivity and the selectivity for the catalytic hydrogenation reaction of nitric acid.
3.2 InJIuence oja catalyst promotor
3.2.1. Measurements with a stationary platinum disc electrode. Addition of germanium dioxide to the sulphuric acid solution (blank), causes typical changes in the voltammogram (Fig. 7). The hydrogen de-
Electrocatalytic hydrogenation processes at controlled potential 1465
x og PTX . 0.5 g Pt/c
E. V vs RHE
Fig. 3. Current-potential relationship for the oxidation of hydrogen at the feeder electrode. Catalyst Pt/C (2Sg/l); electrolyte: 0.1 M KHSO,; o = 2000rev/min.
sorption peaks C and D become smaller, just as the chosen more negative than the potential whereby peak
hydrogen adsorption peak B. The second hydrogen A appears, the height of peak E decreases with
adsorption peak, E, however, becomes larger and a increasing concentration of GeO, in the same manner new anodic peak, A, appears near the oxygen adsor- as the peaks B, C and D. Also, when the start potential ption peak. It was found that the heights of the peeks was chosen near the potential of the oxygen evolution depended on the concentration of GeO,. For the two (anodic start potential), instead near the hydrogen growing peaks (A, E) this is shown in Fig. 8. Also, when evolution, the changes in the cyclovoltammogram
the potential scan rate was varied, the peak currents of were markedly less in the nitrogen saturated solution the peaks A and E varied proportionally to the square than in the hydrogen saturated solution. This can be
root of the scan rate, v. The height of the peak currents explained, assuming that the added germanium dio-
of peaks B, C and D, however, varied linearly with the xide is reduced to germanium at the platinum surface
scan rate. When the anodic scan reverse potential is by means of hydrogen.
.---•-‘--’
x og pt/c
l 059 pt/c
E-02Vvs RHE
1466 J. F. VAN DER PLAS AND E. BAREN~RECHT
Fig. 5. Dependence of the catalyst potential on the rotation speed of the stirrer. Electrolyte: 0.5 M H,SO,,025 MHN03
and 10 - 5 M GeO, ; Catalyst Pt/C (1 g/l).
GeOl + 2H,zGel+ 2H,O. (1)
This reaction can take place in a hydrogen saturated solution and also at the cathodic start potential because of the adsorbed hydrogen. In the nitrogen saturated solution and with an anodic start potential, this reduction process cannot take place. It can only occur during the time interval that the potential of the platinum electrode is in the hydrogen adsorption region. So, only after about six potential scans, notice- able changes in the voltammogram were detected. The decrease in height of peaks B, C and D on the addition of GeOz indicates a decrease in hydrogen adsorption on the electrode surface, which can also be explained as effected by deposition of germanium on the surface. The appearance of peak A must be due to the
a5- 0.4 - 2 03- r > 02- J 0.1 -
redissolving of germanium because of its dependency on the concentration of germanium dioxide, So, the increase of peak E is then the result of the combined effects of decreasing hydrogen adsorption and increas- ing reduction of a germanium species.
Because this peak only increases ifpeak A is present, the reduction process is not the reduction of the Ge02 added, but of a germanium entity formed by redissolv- ing of deposited germanium. The germanium entity formed is apparently Ge ‘* because of the redox potential[ 151 (Table 1) of the reaction
GeF?Ge2+ +2e-. (2)
3.2.2. Rotating platinum disc electrode. When 10 m]
of the germanium dioxide solution is added to a hydrogen saturated sulphuric acid solution, the oxid- ation wave of hydrogen at a rotating platinum disc electrode decreases and even disappears completely when the electrode is treated cathodically in the solution (Fig. 9).
Evidently, the presence of germanium on the elec- trode inhibits the formation of adsorbed hydrogen. When the electrode is scanned continuously after the cathodic treatment, the hydrogen wave returns and its height increases gradually after a number of scans. The dissolution ofgermanium in theanodicpart of the scan obviously cannot be compensated by the deposition of germanium in the cathodic part, due to the fact that the dissolved Ge2+ is swept away from the electrode, so that deposition of germanium is only possible via reaction (1).
3.23. Reduction of nitric acid. Voltammograms have been recorded at a stationary platinum electrode of hydrogen saturated, 1.25 M sulphuric acid solutions containing nitric acid or germanium dioxide and nitric acid (Fig. 10). The addition of nitric acid acid gives only rise to an extra anodic peak at 0.75 V (rhe). As mentioned before, the addition of GeO, changes the blank voltammogram in a much more complex way.
A Blank
0 With Ge02
x With HNO,
l With HNO,+ Ge02
l
Y% \p._ x \ \ rFig. 6. Dependence of the potential of the catalyst on the partial hydrogen Pressure in different solutions, Catalyst: pt/c (1 g/l).
Electrocatalytic hydrogenation processes at controlled potential 1467
Fig. 7. Cyclovoltammogram of added GeO, at a platinum disc electrode. (- - - -) 1.25 M H,S04, H,- saturated (blank) (- ) 1.25 M HZSO+, Hz-saturated with resp. Km5 M GeO, (I), 4.10-s M GeO, (II),
10m4MGeOz (III). u = 1OOmVis.
Addition of both nitric acid and germanium dioxide changes the height of the hydrogen adsorption and desorption peaks still more drastically. Because the adsorption peaks become larger and the desorption peaks smaller, it is evident that the reduction of nitric acid is made possible by the presence of germanium on the electrode surface. When the amount of added germanium dioxide becomes Iarger than 10m4 M, the reduction of nitric acid becomes inhibited. So the effect of the promotor has its limits within a certain con- centration range. For the GeO, promotor we found it to be effective in the range 10m6 - 10m4 M.
4. DISCUSSION
The use of a feeder electrode to influence the potential of the catalyst particles was not effective. although it is evident from our experiments, that a
45 x/“--x x Arwdic peak / x o GMwdic pk 30 - x 2. 67 / / ,/-’ G d.Geo,edd&
Fig. 8. Dependence of the anodic peak A and the cathodic peak E on the concentration of GeO,.
considerable charge transfer takes place between the feeder electrode and the catalyst particles. Due to the relatively high reaction rate of hydrogen at a platinum surface, the main amount of charge transferred be- tween the feeder electrode and a catalyst particle is instantly consumed for the oxidation of hydrogen at the catalyst surface. The charge transferred is then not used to alter the potential of the catalyst particle in order to improve selectivity.
To explain the shift of the potential as a result of a lowering of the hydrogen pressure, the following overall oxidation- reduction reactions on the catalyst surface must be considered :
Electrochemical
oxidation : H,d2H* + 2e- (3)
reduction: 2H+ + 2e-~H, (4)
HNOz + 4H+ + 4e- -+NH,OH + H,O. (5)
Catalytic
HNOS 1 HNO,,, (6)
H,z2H,,
HNOXad + 2 H,, 4 HNO, + H,O. (8) As reaction (6) is very slow, the potential of the catalyst will depend only on the reactions (3) and (4), the reversible Hz/H +-redox reaction. If the pan is lowered, the potential will become more positive according to Nernst, and the rate of the reactions (4) and (5) will decrease. The concentration of nitrous acid therefore increases. Consequently, the redox system of the species H, and HNO, will prescribe a redox potential
different from that of the Hz/I-I-system. Because of
the standard redox potentials of these species (EL, =
0 I-‘, respectively Eho2 = 0.93 V), the resulting mixed potential will be more positive than that of the Hz/H+ system.
1468 J. F. VAN DER PLAS AND E. BARENDRECHT Table 1. Redox potentials of some germanium compounds
Reactm
GeOa + 4H’ + 4C z Get t 2H20
GeO, + 2H+ + 2Q _ GeO + W,O
H,GeO, + 4H+ + 26’ _ -Ge*++ 3HO i:
Ea = -0246
E” = -0206
E” = -0363 + 00295 log [HZGd)SI
CGe”1
li+eo3 + 4H_ + 46 _ Get + 3H,O E” = -0 182 + 0.0148 Log [Hz GeO,]
I 1
Fig. 9. Current-potential relationship of the oxidation of hydrogen at a rotating platinum disc electrode in a 0.5 M H$O., solution. v = lOmV/s. w = 2OOOrev/min.
30 -30 -60 -90 ---- Blank -.- With HNO,
- With HNO, and &Or
Fig. 10. Cyclovoltammogram of a soiution of 25.10-‘MHNOS in O.SMH,SO, with and without the addition of lo-’ M GeO, at a platinum disc electrode. v = lOOmV/s.
Electrocatalytic hydrogenation processes at controlled potential 1469 The addition of Ge03 has evidently a promoting resulting in a change of the catalyst potential due to a effect on the reduction of nitric acid (8). From the shift in the mixed potential of the HJHNO, elec- cyclovoltammetric experiments it follows that an trochemical system.
increase in coverage with germanium on the electrode
surface coincides with a decrease in hydrogen adsorp- Acknowledgement-This work has been carried out with tion. Addition of IO-” M GeOz inhibits the hydrogen financial support from the Netherlands Organisation for the adsorption completely and then inhibits the reduction Advancement of Pure Research (ZWO).
ofnitric acid[3]. So, both hydrogen and germanium on
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5. CONCLUStONS
Charge transfer between a working (feeder) elec- trode and catalyst particles takes place during the oxidation of hydrogen in a slurry cell. However, controlling the catalyst potential via the working electrode is not possible due to the reactivity of hydrogen at the (platinum) surface of the catalyst.
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