Answer:

Explanation:

weak acid (represented here as HA) is one in which the reaction

HA⇌A–+H+(13.3.1)

is incomplete. This means that if we add 1 mole of the PURE acid HA to water and make the total volume 1 L, the equilibrium concentration of the conjugate base A– will be smaller (OFTEN much smaller) than 1 M/L, while that of undissociated HA will be only slightly less than 1 M/L. Equation 13.3.1 tells us that dissociation of a weak acid HA in pure water yields identical concentrations of its conjugate species. Let us represent these concentrations by x. Then, in our "1 M " solution, the concentration of each species is as shown here:

eq3.1-2a.png(1-2)

When dealing with problems involving acids or bases, BEAR in mind that when we speak of "the concentration", we usually mean the nominal or analytical concentration which is commonly denoted by Ca. For example, for a solution made by combining 0.10 mol of pure formic acid HCOOH with sufficient water to make up a volume of 1.0 L, Ca = 0.10 M. However, we know that the concentration of the actual species [HCOOH] will be smaller the 0.10 M because some it ends up as the formate ion HCOO–. It will, of course, always be the case that the sum

[HCOOH]+[HCOO–]=Ca(13.3.2)

For the general case of an acid HA, we can write a mass balance equation

Ca=[HA]+[A–](13.3.3)

which reminds us the "A" PART of the acid must always be somewhere!

For a strong acid such as hydrochloric, its total dissociation means that [HCl] = 0, so the mass balance relationship in Equation 13.3.3 reduces to the trivial expression Ca = [Cl-]. Any acid for which [HA] > 0 is by definition a weak acid.