# Conversion of CO2 to methane

Suppose one could generate an enzyme with arbitrary conditions at the active site. What conditions would then be necessary to cause carbon from dissolved $\ce{CO2}$ in low $pH$ water to combine with the free hydrogen and make methane?

• It's difficult to answer this question; what you're asking about appears to be both purely hypothetical and vague. It would be helpful, in my opinion, if you could be very specific about what you mean by "conditions [...] necessary" to effect the reaction. Also, are you interested in using the water itself as the source of hydrogen? If so, keep in mind that such a reaction requires considerable input of energy (it's typically done by electrolysis) and will not simply occur spontaneously. Enzymes are catalytic, so they will not help in overcoming thermodynamic barriers. – Greg E. Jul 11 '14 at 21:41
• The definition of pH is the amount of free hydrogen in the solution relative to the OH (at least that's how I understand it, correct me if I'm wrong). Therefore water with a high pH would have lots of free hydrogen that could be used to make methane. The only energy input would be adding base to keep the pH constant as hydrogen is used to make methane. – René Jul 12 '14 at 3:00
• The common definition is $pH = -log[H^{+}]$, i.e., colog of hydrogen ion concentration. Hence low $pH$, not high, is associated with high $\ce{H+}$ concentration. Also, "free hydrogen" is not a term I've seen used in reference to $\ce{H+}$, but I have seen it used to describe both neutral molecular and atomic hydrogen. Hence my assumption that you'd be using water as a source for molecular $\ce{H2}$. Also, one most certainly does need energy input. Every reaction has an associated activation energy barrier, at a bare minimum. – Greg E. Jul 12 '14 at 3:21
• The problem with such a reaction is that it does not proceed easily because the activation energy ($E_a$) barrier is high. Methane can be produced from carbon dioxide and molecular hydrogen via, e.g., the Sabatier reaction, but this requires catalysis and high temperatures and pressures to overcome the $E_a$. If you want to extract molecular hydrogen from water, the standard and efficient approach is electrolysis, which yields $\ce{H2}$ and $\ce{O2}$. That reaction, however, is very thermodynamically unfavorable due to the exothermic formation of $\ce{H2O}$, so electrolysis is required. – Greg E. Jul 12 '14 at 3:30
• The reaction does not occur easily. There is no known catalyst to allow reaction to occur at reasonable temperatures in one step. There ARE microorganisms using the reaction you consider as energy source, but the process is not understood well enough and clearly occurs via multiple stages, likely with staged electron transfers toward reducing carbon. – permeakra Jul 12 '14 at 9:43

Let's take a close look at what the reaction entails. You want to convert $$\ce{CO_2}$$ to $$\ce{CH_4}$$ using $$\ce{H^+}$$. It is possible that the hydrogen ion will co-ordinate with $$\ce{CO_2}$$ to produce $$\ce{[CO_2H]^+}$$.
Unfortunately it is very unlikely to have a second hydrogen ion attaching to the same oxygen to produce $$\ce{[CO_2H_2]^{2+}}$$. Now for the impossible step of water leaving and producing $$\ce{[CO]^{2+}}$$. This step would require a large amount of energy and $$\ce{[CO]^{2+}}$$ would be extremely reactive. Since it is highly charged, you would not have any hydrogen ions approaching the moiety.
At this point, any further progress towards $$\ce{CH_4}$$ is fantasy. To stabilize the reaction, you need to add electrons to the carbon (which you don't have a source).
• @Dissenter Changing the pH and using hydrogen peroxide? What will most like be the product of formaldehyde ($CH_2O$) and hydrogen peroxide? I'm guessing a spontaneous reaction producing $H_2$ and $CO_2$. – LDC3 Jul 12 '14 at 13:49