Two Pathways for Partial Methane Oxidation

There are many ingenious methods people have come up with to carry out partial methane oxidation. For the most part, these methods rely on one of two pathways.

Direct Synthesis:

The first of these is to find clever ways to carry out the exact difficult process I described previously: abstract a single hydrogen, and replace it with a single oxygen atom, then stop the reaction.

Many ingenious methods have been devised to make this process more selective and efficient. One popular route is to use single-atom catalytic sites that are only able to remove single hydrogen before they become saturated. Another is to use specially engineered frameworks (known as zeolites) to act as artificial enzymes. These routes show promise but have their own drawbacks, such as being very difficult to synthesize at scale.

Syngas Pathway:

The second pathway takes place in two steps. While it isn’t as efficient as the first pathway, it requires a lot less precision. First, the methane is oxidized under low-oxygen conditions. This removes all the hydrogen, but rather than resulting in water and carbon dioxide, it produces a mixture of hydrogen and carbon monoxide (or, depending on the specific conditions, a mixture of hydrogen and carbon dioxide). This mixture is commonly referred to as synthetic gas, or “syngas”, for short. In the next step, the two components of syngas are reacted together under different conditions to yield methanol.

While this second pathway is quite promising, it is made significantly more difficult by virtue of the two steps of the reaction occurring under totally different conditions. Whereas the first step only takes place at low pressure and low temperature, the second step takes place at high pressure and high temperature. This also presents an opportunity; because the first step is highly exothermic, and produces more gas than it starts with (the number of moles increases), it should theoretically be possible to use the energy released in the first step to power the second step. The question is how to make this work in practice.

The strategy we’re investigating right now is the construction of “nanoreactors”, tiny spheres of metal oxides filled with even smaller particles of metals known to catalyst different the steps of the syngas pathway. The hope is that by confining methane within these nano-reactors, it may be possible to maintain some of the heat and pressure created in the first reaction step to drive the second step.

There are many techniques we will need to develop to make these reactors, and it is unclear how strong an effect it will be possible to generate. However, even if this goal proves impractical, we also hope to use the controlled conditions of the nanoreactor to study the dynamics of the partial methane oxidation reaction and understand how to build more sinter resistant catalysts.

Nanoreactor Design


In this post, I want to describe the three ideas we’re investigating to synthesize nanoreactors right now and the general methodology of nanoreactor design.

The key characteristics of a nanoreactor are:

  1. A closed vessel that can confine metal nanoparticles
  2. Allows limited diffusion of gas through their walls
  3. A scale on the order of tens to thousands of nanometers
  4. Robust enough to withstand harsh reaction conditions
  5. Uniform geometry – all reactors should be a similar size and shape

The first three requirements are necessary for the dynamics we aim to create within the reactors. The fourth requirement ensures that the catalysts we make will not degrade as we perform tests, and will be useful under industrial conditions. The last requirement is not as important but allows us to more easily characterize the performance of nanoreactors, understand them mechanistically, and ensure that multiple batches of nanoreactors will behave similarly.

Mesoporous metal oxides satisfy the first, second, and fourth requirements right off the bat. They are fairly strong materials, resistant to both physical stress and high temperature, are frequently used as inert supports for metal catalysts, and have a complex network of interconnected pores that allow gasses to slowly diffuse through them.


The question then is how to manufacture metal oxide nano-geometries consistently. There is a well-known method to produce uniformly sized beads from silicon dioxide, known as the Stöber process. If a solution of a silicon-containing molecule (tetraethylorthosilicate) is reacted with a solution containing water under basic conditions, it forms silicon dioxide. If this hydrolysis is performed in a well-mixed solution of carefully balanced ethanol and water, it is possible to produce silica spheres of a controlled size, typically on the scale of hundreds of nanometers.


From here, the final challenge is to produce a cavity and insert the catalyst particles that we intend to use. While there are some methods to selectively etch metal oxides, it is only possible under very tightly controlled conditions, and even then is often only possible when the silica is crystalline, rather than amorphous (as is the result of this process). Instead, it is easier to build a new layer on top of the silica sphere, then dissolve the sphere.


This is the general approach we are taking: by adhering the catalyst particles to the surface of the silica, then using a similar approach of hydrolyzing an organo-metallic precursor molecule, we can form a thin film of zirconium or titanium oxide on the surface of the silica spheres, coating the spheres and catalyst. We can then use sodium hydroxide to remove the silica from the inside. Right now, we’re just beginning to test how reliable this coating process is, and whether it is possible to control the thickness, porosity, and other properties of the coating.

Introduction to Partial Methane Oxidation

I just started my internship in the Carngello chemical engineering lab today. The project I’ll be helping with (at least, to start) is designing catalysts for “Partial Methane Oxidation” or PMO. The PMO process is used to convert methane gas (CH4, commonly referred to as “natural gas”) into methanol (CH3OH, also known as wood alcohol).


Methane has found widespread use as a heating gas, but it has two properties that make it unfavourable. The first is that, as a gas with a very low boiling point, it takes a great deal of energy to transport it. To transport it as a gas requires very large vessels, and to hold any appreciable quantity of methane also requires the gas to be pressurized. The more you pressurize it, the more you can store in a container of the same volume (up until it becomes a liquid); however, this also increases the amount of energy you need to expend (pressurization can take up to xx% of the energy you get from burning the methane) and requires thicker containers to hold it (not to mention an increased risk of explosion). In the end, you get ships that look like this designed for transporting methane:


Image result for liquid natural gas ship

(Credit: Fortune Magazine)


This is fine on an industrial scale, but the high infrastructure costs mean that methane extraction is only practical on a large scale. Often, methane occurs in smaller deposits and is accidentally discovered when drilling for oil or minerals. Because of the high cost of liquefying methane, much of it is flared or released, contributing to climate change. Moreover, even when it is being liquefied, some escapes through small leaks in the system.


The other problem that caps methane’s potential is the difficulty of turning it into other chemicals. While we often think of methane as relative because it can be burned, it takes a great amount of energy to start the oxidation process and keep it going. This is part of the reason why methane only combusts at very specific ratios of methane: oxygen.



Converting methane into methanol (a substantially more reactive liquid), solves both of these issues. In order to convert methane into methanol – either for shipping or as a stepping stone towards other industrially useful products – we need to break one of the high-strength carbon-hydrogen bonds and replace it with oxygen.


This sounds simple enough, and it is (it’s the same first step as burning methane). The tricky part is stopping the reaction from going too far, and turning the methane into carbon dioxide.


This is already a challenging task, but the chemistry of methane makes it even more difficult. The first C-H bond is the hardest to break (converting CH4 to CH3), and each successive bond becomes easier to pull apart (So CH3 à CH2 à CH à C). This means that if you have a catalyst that can efficiently “abstract” the first hydrogen, that catalyst is likely to pull off all the other hydrogens as well, which is exactly what we’re trying to avoid.