Hydrogen Peroxide: Why is it unstable?

A while back, I received a lot of good info from a question about hydrogen peroxide.

Followup question: What is it that makes it unstable; i.e., what makes the molecules (presumably caused by the extra hydrogen atom) wanna break up in the first place? I get the idea about how all the molecular commotion generates a lot of heat.

The formula for hydrogen peroxide is H[sub]2[/sub]O[sub]2[/sub], so compared to water (H[sub]2[/sub]O), there is an extra oyxgen atom in the molecule. The oxygen atoms in hydrogen peroxide are not in the most stable state. They have a tendency to want to get to the most stable state, just like a bowling ball at the top of an incline has a tendency to roll downhill.

In the case of the oxygen atoms in hydrogen peroxide, they have an oxidation number of -1. This is not the most stable oxidation state for oxygen. The most stable oxidation number for oxygen is -2 (as is the case for the oxygen atom in water), so the oxygen atoms tend to want to gain electrons to reach this oxidation number, which is the most stable state for room-temperature oxygen.

In the process of gaining electrons, the oxygen in hydrogen peroxide is reduced, and the way it gains electrons is by oxidizing something else, like your hand if you spill it on yourself. :dubious:

Oxygen has very low energy atomic orbitals. That is to say, it is a very electronegative element. In laymans terms, it wants more electrons. You have a simple molecule where two very electronegative atoms (the two oxygens) are deparately fighting over the electrons they have between them. This is one way that hydrogen peroxide can oxidize (steel electrons from) other molecules.

As strange as it may seem, another place it can get these electrons is from another molecule of hydrogen peroxide. When it does this the following reaction occurs.

2H[sub]2[/sub]O[sub]2[/sub] --> 2H[sub]2[/sub]O + O[sub]2[/sub]

Since oxygen is a gas it escapes and the reaction is entropically pushed forward. Of course oxygen itself is pretty electron deficient so it will also steal electrons from other molecules.
Gotta go, I’m sure others will fill in the details I haven’t had time to think through.

No way. Quantum mechanics is in grad school, and it’s staying there. You get electron-deficient sp2 hybridized orbitals that are LOOKING for trouble. Go die now for even reminding me of this.

Then you probably don’t want me to point out that the hybridization of oxygen in hydrogen peroxide is nominally sp[sub]3[/sub]. Of course since each oxygen atom has C[sub]s[/sub] symmetry, the orbitals will be divided up such that none are purely sp[sub]3[/sub]. Perhaps we can call each orbital a linear combination of sp[sub]3[/sub] and sp[sub]2[/sub] hybridization and call it even.

I know way too much QM for a synthetic chemist. Honestly, I wasn’t calling for more QM analysis of the OP in my previous post, I just posted quickly so assumed that I had made some confusing statements that someone else would straighten out. My internet access is a bit sporadic at the moment so I can’t come back and clear up the confusion.

Of this there is no longer any doubt.

Hasn’t anyone told you all you’re supposed to do is shove electrons around?