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In the redox titration of iron(III) with permanganate or dichromate, we use phosphoric(V) acid to "mask" the color of iron(III) because it interferes with the end point color change.

What's the chemistry behind this?

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When phosphoric acid is added to an aqueous solution of iron (III), the ligands in the yellow complexes $\ce{[Fe(OH)(H2O)5]^2+}$ and $\ce{[Fe(OH)2(H2O)4]+}$ ( $\ce{[Fe(H2O)6]^3+}$ are weakly colored but fairly acidic), are successively substituted by $\ce{H2PO4^-}$ and $\ce{HPO4^2-}$. This yields colorless hydrogenphosphatoiron(III) complexes, like $\ce{[Fe(H2PO4)(H2O)5]^2+}$ and the water-free $\ce{Fe(H2PO4)3}$ complex, depending on the phosphoric acid concentration (source). In less acidic solutions, where $\ce{PO4^3-}$ is present, the colorless bis(phosphato)ferrate(III) $\ce{[Fe(PO4)2]^3-}$ is formed (source).

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