How does acidity affect electrode potentials? What is the role of acidity in the decrease of formal potential of $\ce{Fe^{3+}$/$Fe^{2+}}$ in a sulfuric acid and phosphoric acid mixture?

I've observed that the formal potential of the $\ce{Fe^{3+}$/$Fe^{2+}}$ electrode in a mixture of sulfuric acid and phosphoric acid, $E_0 = \pu{+0.61V}$ is much lower than the standard electrode potential, $E_0 = \pu{+0.76V}$.

I came up with a plausible explanation that phosphoric acid would cause $\ce{Fe^{3+}}$ to get precipitated as  $\ce{FePO4}$ and the presence of excess acid will help in the formation of a complex $\ce{[FeHPO4]+}$. Am I correct in my observation? How would the situation change if the medium was basic?
There are numerous reactions that involve iron and many involve either $\ce{H+}$ and $\ce{OH-}$. Take a look at this table and search for "$\ce{Fe}$". The non-standard potential for each one depends on the concentration of the species involved, and for the ones with hydrogen or hydroxide ions this means it includes the pH. This is described by the Nernst Equation. The potential varies as the natural log of the concentration, so you would add or subtract the pH or pOH, depending on the half reaction in question.