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I recently added a potassium hydroxide, $\ce{KOH}$, pellet to a small droplet of hydrogen peroxide (30 %), $\ce{H2O2}$. You could clearly see the potassium hydroxide dissolving in the droplet when instantaneously the solutions began to fizz strongly (boil and/or release oxygen gas, I'm not sure) and release a lot of heat.

I'm not sure what happened, but I've got a few ideas of what might have happened:

  • There is some kind of reaction between the $\ce{KOH}$ and $\ce{H2O2}$.
  • The $\ce{KOH}$ released a lot of energy when dissociating into the aqueous solution and then boiled the water, maybe reaching the temperature of 150.2 °C in which it also triggered the decomposition of $\ce{H2O2}$ into water and oxygen gas.

  • On Wikipedia I read:

    The rate of decomposition increases with rising temperature, concentration and pH.

    So when $\ce{KOH}$ is dissolved, the solution has a high pH and this might trigger the decomposition of hydrogen peroxide

I did notice after a pH test that the plastic lid on which I had condensed some of the evaporated water had a very high pH, so this would most likely contain some dissolved potassium hydroxide.

In case a reaction happened, what could that specific reaction be?

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    $\begingroup$ I have no time to answer, but I am pretty sure that $\ce{KOH}$ is capable of abstracting a proton from hydrogen peroxide, leaving protonated dioxygen ($\ce{HO2-}$), which should be a very strong acid. $\endgroup$
    – Martin - マーチン ♦
    Commented May 5, 2014 at 14:59
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    $\begingroup$ @Martin While $\ce{KOH}$ can certainly deprotonate $\ce{H2O2}$, the resulting hydroperoxide anion $\ce{HO_2^{-}}$ isn't a strong acid. $\ce{H2O2}$ has a $pK_a$ of 11.75 in water according to Wikipedia, and the second ionization is probably several orders of magnitude harder. $\endgroup$
    – Nicolau Saker Neto
    Commented May 5, 2014 at 23:35
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    $\begingroup$ @Nicolau You are very right, I confused it with $\ce{HO2+}$ - as I said I was in a hurry. $\endgroup$
    – Martin - マーチン ♦
    Commented May 6, 2014 at 15:20
  • $\begingroup$ your cap is basic because of small aerosolized water droplets, not some gas phase product. $\endgroup$
    – A.K.
    Commented Dec 13, 2015 at 14:23

3 Answers 3

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According to THE HOMOGENEOUS BASE-CATALYZED DECOMPOSITION OF HYDROGEN PEROXIDE J. Phys. Chem., 1961, vol. 65 , pages 304–306:

$\ce{H2O2 + OH- <=> HO2- + H2O}$

$\ce{H2O2 + HO2- ->}$ [6-membered ring with 2 "H"s bridging 2 "O-O"s, with another H bonded to one of the "O"s] $\ce{-> H2O + O2 + OH-}$

It is clear that the rate will go through a maximum when the peroxide is 50% ionized, that is $[\ce{H2O2}] = [\ce{HO2-}]$.

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  • There is some kind of reaction between the $\ce{KOH}$ and $\ce{H_2O_2}$

There was definitely a reaction.

  • The $\ce{KOH}$ released a lot of energy when dissociating into the aqueous solution and then boiled the water, maybe reaching the temperature of 150.2 °C in which it also triggered the decomposition of $\ce{H_2O_2}$ into water and oxygen gas.

Water cannot reach a temperature above 100 °C unless it is pressurized.

Yes, $\ce{H_2O_2}$ can disassociate rather quickly under the right conditions. An enzyme called catalase, can break up $\ce{H_2O_2}$ at an astounding rate. If you add an amount you can barely see, like the size of a period at the end of a sentence, to 1 mL of 30% $\ce{H_2O_2}$ in a test tube, the oxygen production will push the liquid out of the tube.

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Hydrogen peroxide is stable in acidic conditions. This fact is compounded by the fact that stabilizers (chelates) are commonly acids which do not chelate ions in their basic forms. Furthermore, potassium is a soft largeish ion which catalyzes the autodecomposition.

Bad idea, don't do it again. Hydrogen peroxide is nasty.

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