The ionization of H3PO4 is a stepwise process, and the equilibrium constants for each step are:
$$H_3PO_4 \rightleftharpoons H^+ + H_2PO_4^-: \ K_{a1} = 7.5 \times 10^{-3}$$
$$H_2PO_4^- \rightleftharpoons H^+ + HPO_4^{2-}: \ K_{a2} = 6.2 \times 10^{-8}$$
$$HPO_4^{2-} \rightleftharpoons H^+ + PO_4^{3-}: \ K_{a3} = 4.8 \times 10^{-13}$$
At pH 7.5, the concentration of H+ ions is 3.16 x 10-8 M. This means that the first ionization step is essentially complete, and the concentration of H2PO4- is approximately equal to the concentration of H3PO4.
The second ionization step is also significant at pH 7.5, but it is not as complete as the first step. The concentration of HPO4^2- is approximately 1000 times less than the concentration of H2PO4-.
The third ionization step is negligible at pH 7.5. The concentration of PO4^3- is approximately 100 million times less than the concentration of HPO4^2-.
Therefore, the predominant species of H3PO4 at pH 7.5 are HPO4^2- and PO4^3-.
