Carbon Monoxide Oxidized by OH- not O2?

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I thought that carbon monoxide in normal room air would be oxidized by the oxygen in the air, and that this would happen fairly quickly.

Instead, I was astonished to read that CO is oxidized by the hydroxyl ion. Indeed, most flammable gases in the atmosphere seem to be consumed by hydroxyl ions, not oxygen itself.

I assume the reaction is
CO + 2OH- -> CO2 + H2O

What's going on? Why doesn't oxygen itself oxidize other gases?

I was also stunned to find that the half life of atmospheric CO is on the order of months, not minutes or even hours. Holy cow! The stuff never goes away! Why is it so slow? Is it that you have to have three rather rare molocules bump into each other simultaneously?
 

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  • #2
alxm
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I thought that carbon monoxide in normal room air would be oxidized by the oxygen in the air, and that this would happen fairly quickly.
Why did you think that? You're constantly surrounded with oxidizable substances that don't spontaneously combust. The obvious conclusion would be that just because something can react with oxygen, doesn't mean it does so at room temperature.
Instead, I was astonished to read that CO is oxidized by the hydroxyl ion. Indeed, most flammable gases in the atmosphere seem to be consumed by hydroxyl ions, not oxygen itself.
Where did you read that? Hydroxyl ions (and ions of any kind) are extremely rare in the gas phase.
I assume the reaction is
CO + 2OH- -> CO2 + H2O
That's an unbalanced reaction; where did the negative charges go? You have two electrons more on the left side of the equation.
What's going on? Why doesn't oxygen itself oxidize other gases?
The same reason any reaction won't occur below a certain temperature; the transition-state barrier is higher than the average thermal energy. In the specific case of molecular oxygen, it's because the reactions are often spin-forbidden and cannot occur without forming an energetic radical intermediate.
I was also stunned to find that the half life of atmospheric CO is on the order of months, not minutes or even hours. Holy cow! The stuff never goes away!
Neither do diamonds, yet they're still thermodynamically unstable at room temperature and pressure. This is basic chemical kinetics - reaction rates are governed by the transition state energies, not those of the products. So I don't really know why you'd be surprised. A lifetime of months is actually very short for a gas-phase reaction at these temperatures. If you keep a CO/air mixture in a closed container at room temperature, it will essentially last indefinitely. The reason why it reacts in the atmosphere is because it's getting bombarded with UV radiation from the sun.
Why is it so slow? Is it that you have to have three rather rare molocules bump into each other simultaneously?
Essentially no reactions occur by having three molecules bump into each other simultaneously.
 
  • #3
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Why did you think that? You're constantly surrounded with oxidizable substances that don't spontaneously combust. The obvious conclusion would be that just because something can react with oxygen, doesn't mean it does so at room temperature.
Because, essentially, when I think of things oxidizing in air, I think of them burning in oxygen. As you point out, not necessarily so.

However, I also usually think of solids and liquids not burning in the air at room temperature, because the constituent molecules are bound to each other, and because of the limited surface area.

Despite that, I do indeed hear of many solids and liquids oxidizing on contact with the air, albeit slowly. I'm thinking lead, aluminum, that sort of thing.

When I think about gases, I think of two molecules bumping into each other and being free to react if their thermal energies are high enough. I had not considered that at room temperature, average thermal energies wouldn't be high enough -- which is silly, gas from a lighter or stove doesn't spontaneously burst into flame.

And yet, I had this notion that such collisions would happen often enough that a combustible gas in a room would be consumed eventually. I was apparently wrong, and I'm please to be corrected.

Where did you read that? Hydroxyl ions (and ions of any kind) are extremely rare in the gas phase.
For instance, http://www.answers.com/topic/atmospheric-chemistry" at the answers.com atmospheric chemistry topic. I saw that in several other places as well, but I've closed all my tabs from that search.

http://en.wikipedia.org/wiki/Hydroxyl_radical" [Broken]says, "hydroxyl radicals are produced..., in atmospheric chemistry, by the reaction of excited atomic oxygen with water."

That's an unbalanced reaction; where did the negative charges go? You have two electrons more on the left side of the equation.
Ah. I misread OH-, for the hydroxyl radical, as OH-, the ion. Pardon. I understand now that hydroxyl is neutral. Is the equation otherwise correct?

The same reason any reaction won't occur below a certain temperature; the transition-state barrier is higher than the average thermal energy. In the specific case of molecular oxygen, it's because the reactions are often spin-forbidden and cannot occur without forming an energetic radical intermediate.
Oookay. I'm now taking that intermediate radical to be the hydroxyl radical, although not directly from O2.

Neither do diamonds, yet they're still thermodynamically unstable at room temperature and pressure.
See my fond superstitions about solids v. gases above.

So I don't really know why you'd be surprised.
Because I didn't know any better, and was ignorantly guessing.

A lifetime of months is actually very short for a gas-phase reaction at these temperatures. If you keep a CO/air mixture in a closed container at room temperature, it will essentially last indefinitely. The reason why it reacts in the atmosphere is because it's getting bombarded with UV radiation from the sun.
Gotcha.

Essentially no reactions occur by having three molecules bump into each other simultaneously.
Okay. What exactly happens at the molecular level? I have a CO molecule and an OH radical, and they bump into each other at sufficient energies to combine, but the quantities are not balanced. What happens? Does the CO grab the O and leave the H floating about?

I really don't know, that's why I'm asking. Thanks for getting me up to speed so far.
 
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  • #4
alxm
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Despite that, I do indeed hear of many solids and liquids oxidizing on contact with the air, albeit slowly. I'm thinking lead, aluminum, that sort of thing.
Quite true. This is actually related to the spin-forbidden-reaction thing I mentioned. Molecular oxygen in its ground state has two unpaired electrons which have spins oriented in the same direction (up-up or down-down), AKA a "triplet state". Most molecules have only paired electrons (a 'singlet' or 'closed shell'), and each pair has one "up" and one "down" electron.

The electrons can't spontaneously flip from up to down, so molecular oxygen can't react directly with a closed-shell molecule to form another closed-shell molecule, because one of the electrons needs to flip its spin. A radical, i.e. a molecule with a single unpaired electron ('doublet'), can however flip its spin freely. So two radicals can react and form a closed-shell molecule, regardless of whether their respective spins were 'up' or 'down'.

So oxygen reacts using radical reactions (this is the case in a flame, for instance; lots of radicals flying around). It's also the case in the atmosphere, where you have UV radiation that causes radical formation (which is for instance how oxygen reacts with itself to form ozone in the ozone layer. It's also how CFCs split apart into chlorine radicals which destroy the ozone)

The fact that oxygen can react semi-directly with metals has to do with one of their unique properties: Metals have more or less free valence electrons in them, which are essentially free to change their spins. So oxygen _can_ react with many metals. (But not always easily even then. It usually helps if there are other compounds present as well)

Ah. I misread OH-, for the hydroxyl radical, as OH-, the ion. Pardon. I understand now that hydroxyl is neutral. Is the equation otherwise correct?
Right, then it's balanced. Although the equation still only describes the overall reaction, not the mechanism. Which might be something like:
H2O + UV light -> H* + OH* (splitting of water into radicals)
CO + OH* --> CO2 + H* (forming a hydrogen radical)

You might also have a whole host of other radical reactions going on:
H* + O2 --> HOO* (superoxide radical formation)
HOO* --> OH* + O
O + O2 -> O3 (ozone formation)
2H* -->H2
2O -> O2

These reaction chains are pretty typical of radicals and are important in atmospheric chemistry. Because radicals are quite reactive, since most stable compounds are closed-shell with an even number of electrons. But that a radical reaction often results in creating a new radical, since there's still an odd number of electrons. Two radicals have to find eachother to form a stable compound, usually.
 
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[informative stuff]
Thank you very much. That clears up most of my confusion.
 

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