I know that it can be explained using the Reduction Potential values. But, can it be explained from their structures or something like that ?
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2$\begingroup$ Redox potentials don't explain anything. By this point main effect of knowing std. potentials, or even knowing they exist, is mostly giving you false impression that there is some objective and universal scale to compare them. $\endgroup$– MithoronCommented Mar 29, 2023 at 14:45
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$\begingroup$ Maybe it is due to the chromium being at oxidation number +$6$ in $\ce{K2Cr2O7}$ and manganese at a higher oxidation number (+$7$ in $\ce{KMnO4}$) $\endgroup$– MauriceCommented Mar 29, 2023 at 16:08
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$\begingroup$ They have similar electronic structures and shielding, permanganate has more protons and lower negative charge hence attracts electrons more strongly and is a stronger oxidant. There are many similar trends thruout the Periodic Table: carbonates -nitrates, sulfates-perchlorates, phosphates-sulfates. $\endgroup$– jimchmstCommented Sep 23 at 20:49
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1 Answer
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Structural comparison
The oxidising agents are ionic compounds comprising complex anions and simple potassium cations; the latter behave mainly as spectator ions and don't really participate in any reactions.
If we compare the structures of the permanganate and dichromate ions, we see much similarity in their structure; these comprise mainly double and single bonds between the transition metal atoms and oxygen atoms (the reality is that alternation of single and double bonds in such systems describe a conjugated system with delocalised $\pi$ electrons, such that the actual structure is a resonance hybrid, but that's beside the point).
An example: oxidation of alcohols to carboxylic acids via permanganate
To illustrate how exactly these oxidising agents work, it is useful to search up relevant mechanisms. Below is just one of them expressed in a simple mechanism.
Here we see that the permanganate ions undergo nucleophilic attack and in the process lose an oxygen atom. This is core to what an oxidising agent does; in order to oxidise another species, it must itself be reduced, and losing an oxygen atom is one of the ways it can do so. This is why species like $\ce{LiAlH4}$ and $\ce{NaBH4}$ work instead as reducing agents; by losing $\ce{H}$ atoms, they are themselves oxidised. (See this source for a couple rules of thumb regarding redox)
Anyways, knowing this fact, we can thus say that the major predictor of oxidising power is the strength of the bonds between the oxygen atoms and the central metal atoms; the stronger the bond, the greater the energy requirement for oxygen to be lost and the weaker the oxidising agent.
A simple way to predict this bond strength is in looking at the electronegativity of the metal atoms, which describes its tendency to attract shared electrons. With reference to Ptable, $\ce{Cr}$ has an electronegativity of $\chi = 1.66$ while $\ce{Mn}$ has $\chi = 1.55$. The lower $\chi$ for $\ce{Mn}$ then describes its weaker bonds with oxygen, and thus its stronger oxidising power.
