Question

which one of the following can be used to determine the ka of a weak acid if the concentration of acid is known?ka of a strong acidkb of any weak basepoh of the conjugate base solutionph of the weak acidquestion 8when can we neglect a change in concentration?when the change is much less than the initial concentrationwhen the change is much more than the initial concentrationwhen k is very largewhen q

Explanation:

Step1: Analyze the first question

For a weak acid \(HA\), the dissociation equation is \(HA
ightleftharpoons H^{+}+A^{-}\), and \(K_{a}=\frac{[H^{+}][A^{-}]}{[HA]}\). If we know the concentration of the weak acid \([HA]\) and measure the \(pH\) of the weak acid solution, we can find \([H^{+}]\) (since \(pH =-\log[H^{+}]\)). For a weak acid, \([H^{+}]\approx[A^{-}]\) (from the dissociation equation). Then we can calculate \(K_{a}\).

• \(K_{a}\) of a strong acid is not relevant as strong acids dissociate completely and the method for weak - acid \(K_{a}\) calculation is different.
• \(K_{b}\) of a weak base is related to \(K_{a}\) of its conjugate acid by \(K_{w}=K_{a}\times K_{b}\), but if we only know \(K_{b}\) of any weak base (not the conjugate base of the given weak acid), it won't help us find the \(K_{a}\) of the given weak acid.
• \(pOH\) of the conjugate base solution is related to \(K_{b}\) of the conjugate base (\(pOH =-\log[OH^{-}]\) and \(K_{b}=\frac{[OH^{-}][HB^{+}]}{[B]}\)), and then to \(K_{a}\) via \(K_{w}=K_{a}\times K_{b}\). But this requires more steps (finding \(K_{b}\) first from \(pOH\) of the conjugate base solution) compared to using \(pH\) of the weak - acid solution directly.

Step2: Analyze the second question

In equilibrium calculations (such as for \(K_{a}\), \(K_{b}\), \(K_{c}\) etc.), we can use the approximation \([A]_{initial}-x\approx[A]_{initial}\) (where \(x\) is the change in concentration) when \(x\ll[A]_{initial}\). Mathematically, if \(K=\frac{x\times x}{[A]_{initial}-x}\), when \(x\ll[A]_{initial}\), \(K\approx\frac{x^{2}}{[A]_{initial}}\).

• When the change in concentration (\(x\)) is much more than the initial concentration (\([A]_{initial}\)), the approximation \([A]_{initial}-x\approx[A]_{initial}\) is not valid.
• When \(K\) is very large, the reaction proceeds almost to completion, and we can't use the simple equilibrium - constant expressions (where we assume a small change for weak - electrolyte dissociations) in the same way as when \(K\) is small.

Answer:

For the first question: pH of the weak acid.
For the second question: when the change is much less than the initial concentration.