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I dissolved calcium acetate in about $\pu{80 mL}$ of water and added $\pu{30 mL}$ of 60% peroxide. I should have obtained calcium peroxide which is insoluble in water, but i saw no precipitate. I read somewhere a synthesis where very dilute peroxide was used, so i later tried the synthesis at a maximal concentration of $\ce{H2O2}$ as 6% and still without result.

Should I have used a different salt of calcium? Is it possible the stabilizers in $\ce{H2O2}$ interfered?

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  • $\begingroup$ According to the English edition of Wikipedia (en.wikipedia.org/wiki/Calcium_peroxide), $\ce{CaO2}$ is described both with the low solubility in water mentioned in your question, as well as decomposing in water (see property box). Perhaps the product decomposed before you were able to isolate it? Side note: 60% $\ce{H2O2}$ is about twice as concentrated typically seen in the chemistry lab (stay safe!). On the other hand, if your group has access to this high concentration, possibly a database like Reaxys listing multiple protocols of synthesis may be accessible for you, too. $\endgroup$
    – Buttonwood
    Aug 13 '20 at 20:38
  • $\begingroup$ See also chemistry.stackexchange.com/questions/119807/… $\endgroup$
    – Poutnik
    Aug 14 '20 at 6:51
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One error is the presence of an acid source which results in moving the below equilibrium (written in the reverse of the formation reaction) to the right which occurs slowly with water for many peroxides:

$\ce{ CaO2 + H2O <=> CaO + H2O2}$

Even if you performed the indicated reaction above with H2O2 acting on CaO, you still should check the pH of the hydrogen peroxide, as it is likely acidic. Apparently, phosphoric acid, for example, is frequently employed stabilizer for H2O2 as neutral to basic H2O2 has a much reduced shelf-life.

Further, any presence of a transition metal (Wikipedia on Magnesium peroxide cites an issue, for example, with iron) could also move the equilibrium to the right. Select transition metal oxides are known to produce a cyclic decomposition reaction with H2O2, again effecting the above equilibrium.

My success making a peroxide also entailed working with cold solutions, albeit, in the presence of a stabilizer (like sodium silicate), one may work at warmer temperatures.

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  • $\begingroup$ This brings up another advantage of using a basic source of the metal. You neutralize the added acid and also precipitate transition metals. Magnesia, lime and lithia all do that. $\endgroup$
    – Oscar Lanzi
    Aug 14 '20 at 23:50
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Wikipedia states that the industrial preparation of calcium peroxide is done with calcium hydroxide, not the acetate or a neutral salt. Compare this choice with those used commercially for magnesium peroxide and lithium peroxide.

You want to do these syntheses with strong bases because the metal peroxide itself is most stable in basic conditions. Neutral or acidic water would promote hydrolysis to hydrogen peroxide and thence decomposition.

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