I cannot seek out the reason why while $\ce{H3PO4}$ (phosphorus oxidation +5) is relatively stable and obtainable but $\ce{H3MnO4}$ (the same oxidation state +5) is rarely mentioned and perhaps unobtainable. $\ce{HMnO4}$ could be found on Wikipedia, so I suppose that the stability would decrease by the following order:

$$\ce{HMnO4 > H2MnO4 > H3MnO4}$$

Here is my perspective: I would link the two compounds $\ce{H3MnO4}$ and $\ce{H3PO4}$ via $\ce{H3AsO4}$. I think the key factor that could explain this is the polarity of $\ce{X-O}$ bond. $\ce{Mn(V)}$ has lower electronegativity than $\ce{As(V)}$ so the $\ce{Mn(V)-O}$ bond is more polar and attracts water to give better hydrolysis. The electronegativity of $\ce{Mn(V) < Mn(VI) < Mn(VII)}$, it would probably be the reason why $\ce{HMnO4}$ is the most stable among acids.

Is there anything wrong with my ideas?

  • $\begingroup$ It's a matter of redox, not some "hydrolysis". $\endgroup$
    – Mithoron
    Commented Dec 8, 2023 at 17:22
  • $\begingroup$ Manganese and Phosphorus have nothing in common. Why would you link the two acids $\ce{H3PO4}$ and $\ce{H3MnO4}$ ? What is the meaning of "link" ? And why link via $\ce{H3AsO4}$ ? $\endgroup$
    – Maurice
    Commented Dec 8, 2023 at 17:32
  • $\begingroup$ I was making observation on the stability of oxoacids whose center atoms are similar in oxidation state, which is +5 in the case of H3PO4 and H3MnO4 $\endgroup$
    – Shira
    Commented Dec 9, 2023 at 9:40

2 Answers 2


One difference between manganese(V) and phosphorus(V) is the former can undergo disproportionation with higher manganese oxidation states available.

Not only $\ce{Mn^VO4^{3-}}$ but also $\ce{Mn^{VI}O4^{2-}}$ are stable only as solid salts with an appropriate counterion such as $\ce{K^+}$, or in highly concentrated alkali solutions. Otherwise they tend to disproportionate forming $\ce{Mn^{VII}O4^-}$ and $\ce{Mn^{IV}O2}$. Thus we never can reach a solution acidic enough for $\ce{H2MnO4}$ or $\ce{H3MnO4}$ to be identified. Phosphorus(V) cannot disproportionate in this manner with no higher oxidation state of phosphorus available by typical chemical processes, so $\ce{H3PO4}$ is more plausible.

The $\ce{MnO4^{3-}}$ ion has been substituted for phosphate in synthetic apatites (Wikipedia, citing [1] and [2]). (Manganese in natural apatites would more likely appear in the +2 oxidation state as a trace substituent for calcium.)


  1. K. Dardenne, D. Vivien, and D. Huguenin (1999): "Color of Mn(V)-substituted apatites A10((B, Mn)O4)6F2, A = Ba, Sr, Ca; B= P, V". Journal of Solid State Chemistry, volume 146, issue 2, pages 464-472. doi:10.1006/jssc.1999.8394

  2. Grisafe, D.A. and Hummel, F.A. (1970): "Pentavalent ion substitutions in the apatite structure, part A: Crystal chemistry". Journal of Solid State Chemistry, volume 2, issue 2, pages 160-166 doi:10.1016/0022-4596(70)90064-2

  • $\begingroup$ Can I ask the reason that Mn(V) and Mn(VI) oxoanion is only stable at highly alkali solutions? I was thinking that Mn(VII) oxoanion is more stable due to the stabilization of negative charge, which could be represented by resonance structures. While Mn(V) and Mn(VI) is less stabilized by the decrease of Mn=O bonds. Is it right to assume so? $\endgroup$
    – Shira
    Commented Dec 9, 2023 at 9:57
  • 1
    $\begingroup$ Wikipedia's article on the Mn(V) species suggesrs that a protonated species is an intermediate species in the disproportionation. Also tge disproportiinations consume protons and thus are less thermidynamically favored in stringly basic solutions. $\endgroup$ Commented Dec 9, 2023 at 10:57

Manganese doesn't like to be in +5 and +6 O.S. and the acid form of hypomanganate and manganate anion cannot exist due to rapid disproportionation. So, they exist as salts of alkali metals. Manganese in +7 O.S. is quite stable and hence permanganic acid can still exist in aqueous solution although it can deproptonate or decompose in solution to manganese dioxide, oxygen, and water, with initially formed manganese dioxide catalyzing further decomposition.

Note that manganic acid can exist in a metastable form during the formation of manganate salts but decompose to permanganic acid as it has been noted in an old book(some of the terms are outdated):

Manganic acid in a state of purity has not as yet been obtained; it is known only in combination with bases. In order to obtain it, a strong base, for example, potash or the nitrates of such bases are exposed to a read heat in conjunction with peroxide of manganese (assume MnO2). Whether the operation be performed with or without the contact of air, a manganate is produced; with the contact of air, the peroxide of manganese absorbs oxygen and transformed into manganic acid; without the contact of air, according to Mitscherlich, the peroxide of manganese is decomposed into manganic acid and deuteroxide of manganese (assume Mn2O3)


The manganates in solid state are so intensely green that they often appear black. Their aqueous solution are intensely green also; manganic acid easily decompose to hypermanganic acid (assume HMnO4). All acids, even weakest, impart immediately an intense red color as the green solution of manganates convert into hypermanganic acid, which is accompanied by the formation of black precipitate of peroxide of manganese, with hypermanganic acid further decomposition (if further heat is applied)

Ref.: CHEMICAL ANALYSIS by H.Rose, W. Tegg Publ., 1848

  • $\begingroup$ "Peroxide of manganese". I guess in 1848 they could not distinguish between peroxides as in MgO2/CaO2 and dioxides as in TiO2/MnO2. $\endgroup$ Commented Dec 13, 2023 at 13:00

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