Different solutions in cyclic voltammetry
Recording a CV assess both so-called half cell reactions.  If balanced individually, some of them only need only the transfer of an electron, e.g. for
$$\ce{Fe^{2+} -&gt; Fe^{3+} + e-}$$
Other half-reactions include either $\ce{H+}$ or $\ce{OH-}$, and their consumption / generation alters the pH value of the solution, e.g.
$$\ce{MnO4^- + H2O + 3 e- -&gt; MnO2 + 4 OH-}$$
or
$$\ce{MnO4^- + 4 H3O+ + 2 e- -&gt; Mn^{2+} + 6 H2O}$$
Since each of them are characterized by a standard electrode potential, altering the pH value of the solution thus may alter the reaction pathway accessible; the reaction of permanganate to manganese dioxide is typical for a reaction in a basic aqueous reaction, the one to yield $\ce{Mn^{2+}}$ typical for an acidic one. Indeed pH electrode is an application of this pH dependence which may be quantified by the Nernst equation.
While the examples shown deal with small inorganic ions, the same principle is valid for recording a CV for organic molecules, too.  For the couple quinone / hydroquinone, the variation of the pH value thus may yield a plot like the following:

(reference)
The changes for the redox couple, recorded in different solutions of different pH value equally were plot with the cell potential in function of said variation of the pH value of the analyte:

Note: compare the half reactions with each other.  If both the number of electrons, as well as the number of protons exchanged are equal, the potential of the cell will not be affected by a change of the pH value.
reference: pH Dependent Redox Couple: An Illustration of the Nernst Equation, Walczak et al., J. Chem. Educ. 1997, 74, 1195-1197; doi 10.1021/ed074p1195