Why is hypophosphorous acid (H3PO2) a strong reducing agent?

Why is [hypophosphorous acid](https://en.wikipedia.org/wiki/Hypophosphorous_acid) $(\ce{H3PO2})$ a strong reducing agent? the oxidation state of phosphorous in $\ce{H3PO2}$ is $+5$, as hydrogen and oxygen both are more electronegative than phosphorous, and phosphorous can never increase its oxidation state to more than $+5$, so how can it be a strong reducing agent?


[![hypophosphorous_acid](https://i.stack.imgur.com/78rkS.png)](https://i.stack.imgur.com/78rkS.png)




If you render the oxidation state of hypophosphorous as +5, then the electrons in the phosphorous-hydrogen bonds are attributed to the hydrogen atoms in those bonds. You can do that, but then the hydrogen atoms involved have oxidation state -1 not +1. So the reducing action of hypophosphorous acid is attributed to the negative oxidation state hydrogens.


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The oxidation state of phosphorus in H3PO4 is indeed +5, but in this case you have only 2 and not 4 oxygens. This means that your P is only +1. The reason it's a reducing is because that P really wants to "become" +5.


Oscar Lanzi in his answer suggests that you can assume the H is H-, but this introduces a complication of this being a hydride, and these are not usually stable in solution. Note that this distinction is artificial. Eventually, this is all covalently bonded and the use oxidation numbers suggests ionic bonding.


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