Which of the following solutions is more conductive:

1) 0.1 M sodium hydroxide and 0.1 M hydrochloric acid mixture (effectively NaCl?)

2) 0.1 M sodium hydroxide and 0.1 M acetic acid mixture (effectively NaCH3COO?)

I know the answer is the first one but I can't explain why. I have been taught (high school chemistry) that conductivity of ions is simply relative to [ions]. If that were true, both solutions would be equally conductive. I know that strictly speaking, based on the initial acids and bases, 1) would be more conductive, but I can't understand this because there should be none of the original reactants present anyways. This came up on a test, and I can't justify 1) being correct based on what I have learned.

  • $\begingroup$ What you have been taught is wrong. Conductivities of different ions at the same concentration vary a great deal. $\endgroup$ – Ivan Neretin Jan 25 '18 at 17:20
  • $\begingroup$ Right, but that's why I am trying to figure out how to explain this. $\endgroup$ – Jordan Jan 28 '18 at 18:36

Conductivity of different ions carrying the same amount of charge can vary as per the comment of Ivan. In this question, acetic acid is a weak acid so the final acetic anion will combine with the hydrogen ion from the autodissociation of water molecule to form acetic acid molecules. This is known as the hydrolysis of $\ce{NaAc}$. Hydrochloric acid, on the other hand, is a strong acid so the hydrolysis won't happen. Hydrolysis is the main reason that $\ce{NaAc}$ solution is basic while $\ce{NaCl}$ is neutral.

There should be none of the original reactants present anyways.

This is one fallacy in your reasoning. There are some $\ce{CH3COOH} \rm(aq)$ formed due to hydrolysis in the final solution since it is weak acid.

To summarize, in the first case the conducting particles are equal amount of $\ce{Na+}$ and $\ce{Cl-}$, while in the latter they are $\ce{Na+}$, $\ce{Ac-}$ and $\ce{OH-}$, with the concentration of the first equal the sum of the other two. This is what I believe you need to know. Beyond that without looking up the conductivity table or analyzing the structures of the ions we cannot deduce whose conductance is higher.

In paricular, given a strong acid $\ce{HA}$ and a weak acid $\ce{HB}$, saying that for the solution of the same concentration, $\ce{NaA}$ will be more conductive than $\ce{NaB}$ is simply wrong and there are counterexamples.

  • $\begingroup$ This is exactly what I was looking for. You cannot make that assumption. I know I did a bad job at wording it, but the bottom line is that it would be wrong to assume the weak acid/weak base soln. is more conductive. And I did in fact find the counter examples that you speak of. I will just have a tough job convincing my teacher given that this appeared on a government issued examination... $\endgroup$ – Jordan Jan 28 '18 at 18:25
  • $\begingroup$ @Jordan I believe it will be a hard job. Showing that you have done some research and giving the counter example will be more helpful than providing a link to my answer. Good luck! $\endgroup$ – Weijun Zhou Jan 28 '18 at 18:27
  • $\begingroup$ Thanks a lot, I will try my best. Its just frustrating to see a question like this issued on an exam that thousands of students have written. If you aren't reading into it, you just see two strong acids/bases vs weak acid/strong base and chose the answer accordingly. When you think about it though, it really counters the logic of all ions conduct equal amounts of electricity (I know this is wrong). Just... disappointing that I have spent days unable to explain one question. $\endgroup$ – Jordan Jan 28 '18 at 18:35
  • $\begingroup$ Yes, and it's frustrating as well to see the question being voted down for reasons I don't know... As long as you know that (in my example above) $\ce{HA}$ will be more conductive than $\ce{HB}$ it's fine I guess. $\endgroup$ – Weijun Zhou Jan 28 '18 at 18:41
  • $\begingroup$ Well, good news. I am going to go with the hydrogen oxalate anion as the basis for my argument. It is far more conductive than chloride, yet is the conjugate base of a weak acid. :) $\endgroup$ – Jordan Jan 29 '18 at 4:23

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