4
$\begingroup$

Since it is formed of a strong base and a weak acid it should be a strong electrolyte but a chemistry teacher said that it is a weak electrolyte.....

$\endgroup$
7
$\begingroup$

Sodium acetate is a strong electrolyte in water. The status of sodium hydroxide as a strong base triumphs; it makes even a weak acid such as acetic acid dissociate essesitally 100% into ions (by stealing the protons). And when something is essentially 100% dissociated into ions, by definition it's a strong electrolyte.

$\endgroup$
5
$\begingroup$

A misconception, I feel, that maybe some other students have made after tackling equilibria for the first time. Let me explain by first giving the equilibrium expressions regarding acetic acid and its conjugate base: \begin{align} \ce{CH3COOH + H2O &<=> CH3COO^- + H3O+}& \mathrm{p}K_\mathrm{a} &= 4.76\\ \ce{CH3COO^- + H2O &<=> CH3COOH + OH-}& \mathrm{p}K_\mathrm{b} &= 9.24 \end{align}

I essentially deduced that since both equilibria constants were quite small that both acetic acid and, hence, sodium acetate would constitute weak electrolytes. I have since found out that though my reasoning was correct for acetic acid it wasn't so for sodium acetate.

The reason why is because, in an acetic acid solution, acetic acid is the sole electrolyte, and, hence, the equilibrium expression dictates its dissociation into ions (since it is low there will be slight dissociation). However, as for the sodium acetate solution, the equilibrium expression considers acetate ions completely dissociative from their cations, i.e. it assumes 100% dissociation of the salt. This is indicative of the sodium acetate constituting a strong electrolyte.

$\endgroup$
-1
$\begingroup$

Strong or weak electrolytes will depend on the hydration enthalpy of products in water. That's why any salt like $\ce{NaCl}$ will completely dissolve in the water, though $\ce{NaCl}$ has high bond energy in powder form. In the above case, the hydration of the $\ce{Na+}$ ion will be so dominating that the salt will dissociate (hydrated) completely despite less hydration enthalpy of the anion.

So, as a rule of thumb strong acid/base salts will have more hydration enthalpy but it is just a correlation and will depend on the solvent. The reason for partial or full dissociation is in hydration enthalpy and entropy change.

$\endgroup$

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.