The concentration of the weak acid HF at equilibrium is equal to the initial concentration of the acid the moles per liter of H+.

The statement seems to be incomplete, but I'll try to provide an answer based on what I understand.

The concentration of a weak acid at equilibrium is not necessarily equal to the initial concentration of the acid. This is because weak acids partially dissociate in solution, meaning that not all of the acid molecules break apart into their constituent ions.

For the weak acid HF, the dissociation in water can be represented by the following equilibrium:

HF ⇌ H+ + F-

The equilibrium concentrations of HF, H+, and F- are related by the acid dissociation constant Ka:

Ka = [H+][F-] / [HF]

If we let x represent the change in concentration of HF due to dissociation, then at equilibrium we have:

[H+] = x [F-] = x [HF] = initial concentration of HF - x

Substituting these into the Ka expression gives:

Ka = x^2 / (initial concentration of HF - x)

This equation can be solved for x, which gives the equilibrium concentrations of H+ and F-. Note that x is not necessarily equal to the initial concentration of HF, especially if the initial concentration is large.