2
$\begingroup$

Question:

Classify each of the following ions according to whether they react with water to give a neutral, acidic, or basic solution:

As a general rule:

Strong acid + Strong base $\rightarrow$ Neutral solution

Strong acid + weak base $\rightarrow$ acidic solution

Weak acid + Strong base $\rightarrow$ Basic solution

According to my solutions manual.

a) $\ce{F-}$ is the conjugate base of a weak acid, therefore the solution is basic.

b) $\ce{Br-}$ is the anion of strong acid, therefore the solution is neutral.

c) $\ce{NH4+}$ is the conjugate acid of a weak base, therefore the solution is acidic

d) $\ce{[K(H2O)6]+}$ is the conjugate base of a weak acid, therefore the soltuion is basic.

e) $\ce{SO3^2-}$ is the conjugate base of a weak acid, therefore the solution is basic

f) $\ce{[Cr(H2O)6]^3+}$ is an acidic solution, therefore the solution is acidic.

However, all these answers imply that water is acting as a strong base. How is this possible when water pH is $\sim 7$?

$\endgroup$
2
$\begingroup$

You have to take into consideration the pKa and pKb, not the pH, when making judgments about the acidity or basicity, respectively, of a molecule. In the examples you pulled out, for water to act as a base, it must have a higher propensity to grab a proton or donate electrons than the competing molecule. Thus, it must have a higher pKa or lower pKb than the other molecule to act as a base.

$\endgroup$
4
$\begingroup$

Question:

Classify each of the following ions according to whether they react with water to give a neutral, acidic, or basic solution:

As a general rule:

Strong acid + Strong base $\rightarrow$ Neutral solution

Strong acid + weak base $\rightarrow$ acidic solution

Weak acid + Strong base $\rightarrow$ Basic solution

That's not true. The only time in which these rules are true is when you have an equimolar concentration of the acid and base in question, and the acids and bases must both be monoprotic and monobasic, or diprotic and dibasic, etc. Otherwise there are too many variables to consider for these guidelines to always ring true. Consider mixing 1.0 molar $\ce{H2SO4}$ with 0.10 molar $\ce{NaOH}$ ... I guarantee you that the pH of that solution is not "neutral."

d) $K(H_{2}O)_{6}^+$ is the conjugae base of a weak acid, therefore the soltuion is basic.

I didn't know that potassium hydrated to a significant extent in water; even so, what the solution manual has written down is not the conjugate base of hydrated potassium cation - the weak acid.

However, all these answers imply that water is acting as a strong base. How is this possible when water pH is $\sim 7$?

How do these answers imply that water is acting as a strong base?

Also regarding how water can be a strong base in general - it just depends on your solvent system.

$\endgroup$
-3
$\begingroup$

One more thing that you have to keep in mind is that water acts as a strong base ONLY with weak acids.

PECULIAR nature of water is that with weak bases, water acts as an acid. So, it can be said that it is amphoteric in this regard.

$\endgroup$
  • $\begingroup$ Water is never a strong base, water is always a weak base. The strengths of acids and bases are intrinsic properties of the corresponding molecules/atoms, only dependent on the solution to a lesser extent. $\endgroup$ – Jan Feb 4 '17 at 15:44

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.