Van't Hoff factor for citric and phosphoric acids
The Van't Hoff factor isn't an intrinsic property of any species - it depends on the dissociation degree of an electrolyte. Both of those acids are weak and triprotic, meaning they can lose up to three protons in solution (depending on the pH). Therefore, the Van't Hoff factor will likewise change depending on the pH.

1) If the acid is the only species in solution, and since it's a weak acid, you need to use the same formula as for any weak electrolyte:

$i=1+\alpha$

You can figure out the dissociation factor, $\alpha$, from the dissociation constant:

$K_1=\cfrac{[H_2A^-][H^+]}{[H_3A]}=\cfrac{(\alpha C_{H_3A})^2}{(1-\alpha)C_{H_3A}}=\cfrac{\alpha^2}{1-\alpha}C_{H_3A}$

where $C_{H_3A}$ is the formal concentration of the acid.

2) If you have more species in solution, it will depend on the pH of the solution. In general, you'll have this expression:

$i=\cfrac{[H_3A]+[H_2A^-]+[HA^{2-}]+[A^{3-}]+[H^+]}{C_{H_3A}}$

After some tedious substitutions you'll get

$i=1+\cfrac{K_1}{[H^+]+K_1}+\cfrac{K_2}{[H^+]+K_2}+\cfrac{K_3}{[H^+]+K_3}$

Some relevant values:

$pH=0 \rightarrow i \approx 1$

$pH=pK_1 \rightarrow i = 1.5$

$pH=pK_2 \rightarrow i = 2.5$

$pH=pK_3 \rightarrow i = 3.5$

$pH=14 \rightarrow i \approx 4$

and it changes logarithmically in between.