How to properly treat potassium-iron oxide catalysts

Over the past few months I tried to collect as much information as I could find on the iron-potassium oxide catalysts routinely used in the work by Holmlid et al. After much research and note taking, it seems apparent to me that the probable main reason why experimenters have been having several difficulties with replicating the results is the preparation involved for activating the catalysts, or in other words to render them catalytically active.

This information is widely published in the mainstream scientific literature, but this is not immediately obvious by just reading Holmlid's papers. In fact, very little is directly available about them just by reading his papers. One has to follow the references and read the related papers cited there. These are the two main works he usually cites when referring to these catalysts:

• G. R. Meima and P. G. Menon, Appl. Catal. A 212 (2001) 239.
• M. Muhler, R. Schlögl and G. Ertl, J. Catal. 138 (1992) 413.

To cut a long explanation short, it seems that upon specific conditions, in particular those similar to the industrial processes they're normally involved with, these catalysts undergo various structural transformations. The main defining characteristic is the formation of an unstable potassium ferrite (KFeO₂) compound on their surface, which is generally considered to be the main catalytically active phase (this compound in isolation is as active as the industrial catalyst).

My understanding is that this compound is formed during these conditions:

• During catalysts synthesis from precursor materials at about 800-850 °C in air
• During actual process conditions in the presence of hydrocarbons and oxygen at about 600-650°C (typical operating temperature)
• Simply by heating the catalyst in a non-reducing residual atmosphere under vacuum conditions above 600°C

The reaction is as follows:

Fe₂O₃(s) + K₂CO₃(s) => 2 KFeO₂(s) + CO₂(g)

Interestingly, the compound is characterized by a dark olive green color, which can be used to detect its formation. Under ambient air conditions the compound is reported to quickly decompose back into red hematite (Fe₂O₃) due to CO₂ absorption, with water catalyzing the reaction. "As received" catalysts from long-term storage under uncontrolled conditions therefore are mostly composed (at least on the surface) of hematite and potassium carbonates, and are not immediately catalytically active. Any activation procedure would have to be preferably done in-situ.

A few weeks ago to test the hypothesis that just calcinating (heating in air) for a relatively short period of time pre-formed, commercially available industrial iron oxide catalysts similar to those used by Holmlid could restore at least part of the potassium ferrite surface layer (by the above reaction) with a possibly noticeable change of color, I asked Alan Smith (our LENR-Forum moderator), who had a certain amount of these catalysts at disposal, to calcinate some of them for 90 minutes at 900°C and take a couple photos before and after the treatment. The entire procedure overall required minimal effort on his part.

It did appear that the "after" catalysts changed slightly in color, at least in part confirming existing information in the literature and that at the right conditions potassium ferrites (KFeO₂) might have been indeed formed. Note that the white patina visible in the "before" photos is likely potassium carbonate. Unfortunately I haven't been provided photos showing their state several hours later after storage in air. From various reports it seems that this occurs visually within hours. The so-treated catalysts have not been tested in an experimental apparatus under a hydrogen atmosphere either.

I'm wondering if other experimenters with access to basic materials and equipment would be interested in performing some tests to confirm some of the characteristics of these catalysts. I can provide all references and information needed, but in absence of public of interest on the subject putting too much effort early on this thread would probably not be very useful.

Potentially interesting activity diagram from http://dx.doi.org/10.1006/jcat.2002.3725

References suggesting that the KFeO₂ compound formed has a dark olive green color. Plausibly, this might be mistaken for a black color or reduction to Fe₃O₄ which is black and non-active. At low concentrations it can give a light ochre or beige coloration (observation from earlier tests at lower—insufficient—temperature with the same catalysts used above) to the rest of the catalytic material, but prolonged heating to elevated temperature also in a vacuum should increase its concentration on the surface due to diffusion (and desorption) processes taking place, eventually rendering the change more apparent.

• https://doi.org/10.1080/01614947408071864 (p. 298; catalyst turns "black" by heating in a vacuum at 600 °C)
• https://doi.org/10.1007/BF00766208 (p. 209)
• https://doi.org/10.1016/0021-9517(90)90003-3 (p. 346)
• https://doi.org/10.1016/0021-9517(92)90052-J (p. 535, 540, 541)
• https://doi.org/10.1016/0021-9517(92)90295-S (p. 436)
• http://lib.dr.iastate.edu/cbe_pubs/222 (p. 2114, 2116, 2119)
• http://lib.dr.iastate.edu/cbe_pubs/232 (p. 7431)
• https://doi.org/10.1006/jcat.1996.0318 (p. 188, 190, 193)
• https://doi.org/10.1002/zaac.19633210505 (p. 249, 259; "olivgrün")
• https://doi.org/10.1016/S0926-860X(02)00270-3 (p. 223)
• https://doi.org/10.1016/0926-860X(95)00314-2 (p. 75, 79; dark green, from K-Mn oxide, partially related)
• https://doi.org/10.1016/0926-860X(96)00173-1 (p. 392, 395, 398)
• https://doi.org/10.1016/j.jssc.2006.01.054 (p. 1427)
• https://doi.org/10.1016/j.apcata.2008.11.007 (p. 51)
• https://doi.org/10.1016/j.apcata.2011.07.036 (p. 102)

Among other things, the references above may contain useful information regarding the procedures used to synthesize KFeO₂. It can be useful to know for example that temperatures significantly greater than 800 °C have to be avoided in order to not synthesize K₂Fe₂₂O₃₄ instead (if the aim is only obtaining KFeO₂), which unlike the other is very easily reduced in a pure hydrogen atmosphere. In actual catalysts however the active state is apparently a metastable balance between the two compounds.

Below are some excerpts in image form suggesting that potassium ferrite can also be formed from "as received" catalysts in the process of heating in a vacuum (due to potassium diffusing from the bulk to the surface of the catalyst and reacting with free Fe₂O₃ phases), and/or that it's unstable.

Source: https://doi.org/10.1080/01614947408071864

Source: http://lib.dr.iastate.edu/cbe_pubs/222 (TGA/Thermogravimetric analysis device had a pressure of 0.003 atm)

Source: https://doi.org/10.1021/ie00023a010 (Holmlid, 1993. Note that I cut much of the original text in the excerpt)

Some might be asking what is there so special about this 600 °C threshold mentioned in the excerpts of the previous comment.

It appears that there's a temperature range (550–650°C) from which the emission of potassium atoms and ions increases to significant levels compared to lower temperatures. Not only this includes the emission of excited (Rydberg) states from the surface, but also the diffusion of ions from the bulk to the upper layers of the catalyst. This means that under active conditions the surface of the catalyst will slowly develop an increased concentration of the active KFeO₂ compound.

A typical desorption signal graph looks like this:

Source: https://doi.org/10.1023/A:1019013504729 (Available on Researchgate)

In a more user-friendly linear axis form it becomes like this, from which it becomes clearer that at lower temperatures there is little hope that structural changes in the catalyst will take place in a reasonable amount of time (since as previously mentioned they are driven by diffusion of potassium to the surface):

Other sources showing similar graphs and behavior from these compounds and industrial iron-potassium oxide catalysts:

• http://dx.doi.org/10.1016/j.jcat.2007.02.009
• http://dx.doi.org/10.1016/j.cattod.2010.03.059
• http://dx.doi.org/10.1039/c5cp04108b
• https://doi.org/10.1142/S1793604711001865
• https://doi.org/10.1016/j.jcat.2006.01.026

It might be interesting to know that at least under a low partial H₂ pressure (in the order of a few millibar) this active KFeO₂ compound will not get easily reduced, not even at 630 °C—the typical operating temperature of the actual industrial catalyst in the styrene process (although apparently reduction will start to take place above 700°C). It seems that once formed it increases the reduction resistance of the industrial iron oxide catalysts, even when present in small amounts, at least under less reducing conditions than a pure H₂ environment.

In the styrene process, which generally operates at higher sub-atmospheric pressures, steam is admitted with ethylbenzene to keep the catalyst in an active oxide form. When using hydrogen only, either a low pressure or low temperatures would have to be used to prevent irreversible reduction processes, taking place with the following reaction as indicated in https://doi.org/10.1021/ie00023a010:

But perhaps if the catalyst was mostly composed of the catalytically active phase(s) this would be less of a concern.

Relevant excerpts from http://dx.doi.org/10.1016/j.apcata.2008.11.007

Other related sources:

• http://dx.doi.org/10.1016/j.apcata.2011.07.036
• https://lib.dr.iastate.edu/cbe_pubs/222 (especially) [open access]
• https://lib.dr.iastate.edu/cbe_pubs/232 [open access]
• https://lib.dr.iastate.edu/rtd/1186 [open access]
• https://doi.org/10.1016/0021-9517(90)90003-3

Catalyst deactivation phenomena in styrene production by Meima and Menon (2001)

This paper, which has been always referenced by Holmlid over the past years in his ultradense hydrogen studies (it's the newer of the first two cited in the opening post in this thread), offers a general overview of K-Fe oxide catalysts for the production of styrene in the context of deactivation processes. Indirectly and upon careful reading it helps understanding what makes them work properly. A few cited patents (in refs 36 and 37) might be of interest.

Some time ago I took notes of what I considered the most interesting bits while reading it. Here they are for everybody's benefit, in case there's anybody interested (not that so far I have seen much signs of interest, though). I can't guarantee that they're accurate, so of course the best thing would be reading the actual paper ^{(e.g. using sci-hub).}

Quick notes

• The styrene process with these catalysts usually runs at 600–650 °C in a partial vacuum.
• It has been observed that removing coke (carbon deposit) formation with steam quickly increases activity. Coke formation however is not always harmful. Menon classifies it into several types [6].
• KFeO₂ is generally considered to be the active phase of these catalysts. First observed by Hirano [4, 14] then other groups.
• The active state of the catalyst is an equilibrium between KFeO₂ and K₂Fe₂₂O₃₄. Reduction by hydrogen causes formation of inactive Fe₃O₄ [19].
• Stobbe [16], Goodman [20], Lundin et al [21], Menon and Holmlid [22] also support the notion of the active state being KFeO₂.
• Holmlid et al [24, 25] observed excited states and clusters (Rydberg matter) from used and fresh catalysts at 600–800 °C in an UHV chamber. An energy diagram has been constructed taking into account the mechanisms proposed earlier by Ertl et al [19].
• In the actual styrene production process the fresh catalyst is partially reduced to Fe₃O₄, but becomes more stable afterwards.
• Muhler et al [19,28] note that the precursor materials contain excess Fe³⁺ ions which form the unstable ternary compound K₂Fe₂₂O₃₄, which forms in turn KFeO₂ which is considered the active phase.
• As the above occurs, Fe₂O₃ is slowly reduced to Fe₃O₄ which is considered inactive.
• K₂Fe₂₂O₃₄ acts as a K storage medium to prevent reduction of Fe³⁺; once depleted, the catalyst becomes irreversibly deactivated.
• In general, it is acknowledged that transformations during actual usage occur.
• Reduction (to Fe₃O₄) of Fe³⁺ -containing phases must be prevented and KFeO₂ possibly works beneficially towards this goal.
• Promoters—including especially Cr—also stabilize the Fe³⁺ oxidation state, i.e. prevent or slow down reduction to Fe₃O₄.
• Reduction to Fe₃O₄ also physically weakens the catalyst, causing a transition from hexagonal lattice to cubic lattice structure.
• The loss of the K promoter increases coke formation, which clogs catalyst pores, decreasing the number of active sites.
• Alkaline earth oxides as promoters tend to improve structural stability. A number of studies and patents exist showing their effect.
• In conclusion, catalyst deactivation processes are complex. They can be of both permanent and temporary nature. Much research work has gone towards decreasing the former.
• Stabilizing the Fe³⁺ oxidation state seems to be the main trend (as of 2001). This also has the effect of increasing the K desorption threshold temperature as proposed by Holmlid et al [33].
• Patents and studies exist on this aspect too. For example, MgO or ZrO₂ appear to be good support materials that can stabilize these catalysts.
• The pore structure of the catalyst is also important on this regard. Notably, treatment in hydrogen under specific conditions can alter it to prevent further deactivation [37]. This might also promote KFeO₂ formation.

Relevant references

• [4] T. Hirano, Bull. Chem. Soc. Jpn. 59 (1986) 2672.
• [6] P.G. Menon, J. Mol. Catal. 59 (1990) 207.
• [14] T. Hirano, Appl. Catal. 26 (1986) 65.
• [16] D.E. Stobbe, F.R. van Buren, A.J. van Dillen, J.W. Geus, J. Catal. 135 (1992) 533 and 548.
• [19] M. Muhler, J. Schutze, M. Wesemann, T. Rayment, A. Dent, R. Shlögl, G. Ertl, J. Catal. 126 (1990) 339.
• [20] K. Coulter, D.W. Goodman, R.G. Moore, Catal. Lett. 31 (1995) 1.
• [21] J. Lundin, L. Holmlid, P.G. Menon, L. Nyborg, Ind. Eng. Chem. Res. 32 (1993) 2500.
• [22] L. Holmlid, P.G. Menon, Appl. Catal. A: Gen. 212 (2001) 247.
• [24] J. Lundin, K. Engvall, L. Holmlid, P.G. Menon, Catal. Lett. 6 (1990) 85.
• [25] K. Engvall, L. Holmlid, P.G. Menon, Appl. Catal. 77 (1991) 235.
• [28] M. Muhler, R. Shlögl, A. Reller, G. Ertl, Catal. Lett. 2 (1989) 201.
• [33] L. Holmlid, K. Engvall, C. Åman, P.G. Menon, in: Guczi, et al. (Eds.), Proceedings of the 10th International Congress on Catalysts, Budapest, 1992, Elsevier, Amsterdam, 1993, p. 795.
• [36] D.M. Hamilton WO/18593 and WO/18594 (1995), to Shell.
• [37] Y.T. Shiraki, J.I. Matsui, US Patent 5,824,831 (1998), to Idemitsu.

magicsound

I can confirm that the test was performed with catalyst pellets provided by you. Thanks to you too for making that test possible.

After extensive reading it's now clearer that the potassium on the surface from the analyses was probably mostly present in the form of carbonates, as its oxide readily absorbs CO₂ from the atmosphere at room temperature. The strongly red appearance of the "as-received" pellets should suggest that the Fe oxide was mostly hematite.

I think a good test requiring the least amount of efforts and resource expenditure while still being useful would be documenting the stability over time of this active KFeO₂ compound, which should be formed in limited amounts on the surface of the catalyst upon prolonged calcination in air at atmospheric pressure at about 800°C by the reaction of Fe₂O₃ with K₂CO₃. Higher temperatures promote the formation of other phases. Shell 105-type catalysts are usually calcined during their manufacturing for about 2-3 hours at temperatures in the 850-950 °C range.

Source: http://dx.doi.org/10.1039/b104263g

How exactly unstable this compound is in a standard atmosphere? It's been reported in the literature to revert to red hematite within a few hours, but I suspect the bulk of the change occurs within the first couple dozen minutes of exposure to ambient air. A timelapse video could be interesting and be something unique as this sort of visual documentation doesn't appear to be available.

Source: http://lib.dr.iastate.edu/cbe_pubs/222 (open access)

More good photos of the freshly calcined pellets showing a more apparent olive green coloration than the one I uploaded would also be good to have. Photographing them together with uncalcined catalyst samples would make any difference more visible.

It's likely that crushing the pellets into sub-millimeter granules (but not a too fine powder) could yield better results. A comparative test with the cylindrical pellets after the same treatment might be of interest. In that case it would be also be interesting to know if such crushed pellets from long-term storage also show the same deep red coloration on the inside (probably they do).

After calcination, would the cylindrical (uncrushed) pellets, once they show a decomposed red appearance, also appear decomposed on the inside? It would be interesting to crush some to find out.

I think these tests could be done in a single run before and after calcination in a furnace at 800°C for several hours. Anything between 4-6 hours at this temperature would be fine; more than this is probably not needed but it will not harm either. In the following study the authors synthesize KFeO₂ directly from powdered precursors with 3.5h under flowing air.

Source: http://lib.dr.iastate.edu/cbe_pubs/232 (open access)

I believe this more or less sums up what can be done with a standard furnace to find out more about the temporary nature of the active phase of these catalysts. It would be great if you could perform some of these tests. Thanks for donating some of your lab time!

On a partially related note, in the industry the actual styrene catalysts are activated simply by exposure to reaction conditions (about 600-650 °C with steam and ethylbenzene at pressures < 1 atm). Steady state conditions are achieved generally within a few days of activity.

Source: https://doi.org/10.1002/9783527610044.hetcat0163

magicsound

That would be a great addition to the other tests. It seems that exposure to most atmospheric compounds (N₂, O₂, CO₂, H₂O) at room temperature make such active phase decompose. Argon should be fine, though.

Source: https://doi.org/10.1002/9783527610044.hetcat0163

Source: https://doi.org/10.1016/0021-9517(90)90003-3

In reference to the instability of potassium ferrite (KFeO₂) also cited in comment #13, Kotarba et al. noted in http://doi.org/10.1023/A:1019013504729 (open access through Researchgate) :

A partial pressure of water > 0.01 atm in the atmosphere appears to negatively affect it. Storage in vacuo might be the preferred long-term solution.

Seiber Troutman

If you want to contribute constructively to this thread you could try synthesizing potassium ferrite (KFeO₂) on your own and note any short-term interesting change in color of the starting hematite.

The resulting material should appear olive green upon synthesis and it's the one that actually gives the slight green tint to the industrial catalysts upon calcination—as well as making them apparently "turn black" upon heating in a vacuum due to emission of excited states, as linked in one of the firsts posts in this thread. It should also be rather unstable under ambient conditions, more than the treated catalyst pellets themselves.

Apparently this compound alone is as catalytically active as the industrial catalysts. So, potentially (there are some caveats) there wouldn't be necessarily need to hunt for hard to obtain catalysts for replications of experiments employing them.

Materials (in equal proportions by weight to simplify things out—in practice usually slightly more potassium carbonate than hematite is used):

• Fe₂O₃ (hematite) powder
• K₂CO₃ (potassium carbonate)
• H₂O

Proposed procedure:

1. Place the solid materials in a crucible and add just enough water to obtain a saturated solution of K₂CO₃. About the same amount of water by weight should be sufficient according to solubility data provided on Wikipedia.
2. Stir well as to obtain a more or less homogeneous dense slurry.
3. Place the crucible in the furnace, increase temperatures to 200°C and hold for 30 minutes to let the solution dry (should be enough).
4. Remove the crucible from the furnace and crush the resulting powder with a pestle to homogenize it.
5. Place again the powder in the furnace, increase temperatures to 800-825°C and let it calcinate for 4–6 hours.
6. Remove the crucible from the furnace and quickly take photos of the calcinated sample.

Source: https://doi.org/10.1016/0021-9517(92)90295-S

Source: http://lib.dr.iastate.edu/cbe_pubs/232

Seiber Troutman

They're not the same thing.

KFeO₂ = potassium ferrite. It's unstable in air and according to the various papers linked so far decomposes quickly into red hematite potassium carbonate, in addition to other iron and potassium compounds. It's reportedly the active phase of the catalysts of this thread.

K₂FeO₄ = potassium ferrate: https://en.wikipedia.org/wiki/Potassium_ferrate

Seiber Troutman

Not quite. Styrene is the precursor compound (a liquid hydrocarbon) used to produce its polymerized form (a solid) called polystyrene, which is what many plastic materials are made of. However, in the context of this thread this is not really important, as using these catalysts for styrene production isn't the end goal here.

Quote from magicsound

The test is now under way. I'll post photos at the conclusion.

Thanks again. Looking forward to reading and seeing the results.