physical chemistry - What does it mean for a reaction to favor the reactants/products?

(This example is purely hypothetical.)

You have the reaction

$$\ce{H2(g) + O2(g) <=> H2O2(g)}$$

at $T = 500\ \mathrm{K}$. The reaction reaches equilibrium at the following concentrations:

$$\ce{[H2]} = \ce{[O2]} = 5 \times 10^{-3}\ \mathrm{mol\ dm^{-3}}$$ $$\ce{[H2O2]} = 4\times10^{-5}\ \mathrm{mol\ dm^{-3}}$$

This gives

$$K_c = \left(\frac{\ce{[H2O2]}}{\ce{[H2][O2]}}\right) = 1.6$$

However, the total concentration of the reactants is $250$ times higher than the concentration of the product.

Still, as $K_c > 1$, per definition, the products are favored.

This seems counterintuitive to me, and my Chemistry teacher couldn't really explain it to me, so I was hoping someone here could explain why, even when the reactants are so much more plentiful, the products are considered to be favored.