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For example, is 3M HCl a stronger acid than 1M HCl?

I would reason that the concentration of an acid/base does not influence its strength. Strength is determined by the pKa, and, as per Le Chatelier's Principle, the initial concentration does not influence the equilibrium constant.

It may shift the equilibrium concentrations (meaning that the pH is higher for the 3M HCl), but it will not change $K_a=\frac{[H_3O^+][Cl^-]}{[HCl]}$ at equilibrium.

Is this right?

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    $\begingroup$ Your reasoning is incorrect. $\endgroup$ – MaxW Dec 5 '18 at 0:38
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    $\begingroup$ What's hard about this? 3 > 1 so there is more acid and hence it is stronger in that sense. Also HCl ionizes fairly completely so the pH is different too. $\endgroup$ – MaxW Dec 5 '18 at 0:46
  • $\begingroup$ @MaxW, What then determines the strength of an acid? Is it the amount of H+ or H3O+ ions in the solution? How is that any different from the pH? $\endgroup$ – Daniel Dec 5 '18 at 0:53
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    $\begingroup$ It means nothing, unless you define it to mean something. That's why, BTW, scientists tend to stick to the terms with precise meaning (concentration, pKa, etc.) $\endgroup$ – Ivan Neretin Dec 5 '18 at 6:19
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    $\begingroup$ Whole solution is an acid - compound may act as acid or not. Calling HCl itself an acid is a sort of classification made with usual function. $\endgroup$ – Mithoron Dec 6 '18 at 0:02
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The term ‘strong acid’ is sometimes used in a rather fuzzy way and you ran into problems doing so. I prefer to use the term ‘strong acid’ only with respect to an acid’s $\mathrm pK_\mathrm{a}$ value and disregard all other influences. This gives a clearly defined measure of acid strength and we can easily sort various acids by their strength into stronger or weaker acids.

However, this is looking at the acid as a molecule. In real-life applications you are typically more interested in the property of a solution. To give a real-world example, imagine a $\pu{1M}\ \ce{HBr}$ solution and a $\pu{12M}\ \ce{HCl}$. solution. Obviously, $\ce{HBr}$ is the stronger acid, but the concentration of $\ce{HCl}$—also a strong acid and thus fully deprotonated—is higher. Therefore, the $\ce{HCl}$ solution is more concentrated or, as some would say, more acidic. It can do greater harm and it is able to protonate more Brønsted base molecules than its $\ce{HBr}$ counterpart.

If instead of examining a $\pu{1M}$ and a $\pu{12M}$ solution I had examined a $\pu{e-3M}$ and a $\pu{1M}$ solution, we could even base the discussion around the solution’s resulting pH: the stronger acid is much more diluted and will result in a solution of pH 3 while the weaker acid results in a solution of pH 0.

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Normally, when comparing Acid Strengths between two acids, the first criterion used is their respective value for Ka at the same temperature. Obviously the one with a higher Ka at the same temperature is "stronger".

If you have two different acids (let's say HF and HCl) with the same pH, their strength can be compared in terms of their conversion ratio (a ratio that represents acid dissociation in solution).

Conversion = ax/CAo

Where:

a = stoichiometry coefficient for the acid in its dissociation reaction;

x = change in acid concentration to reach equilibrium;

CAo = Initial concentration of acid

For both HCl and HF, a = 1 due to their corresponding dissociation reactions.

For HCl, it can be assumed that CAo = x, so its Conversion is 1 or 100%

For HF, CAo >> x, so its Conversion is << 1 or << 100%

Therefore, HCl is the stronger acid.

However, since you're comparing the same acid (HCl) at different concentrations (and same Temperature, I'm assuming), concentration should be used as a secondary criterion because both the Conversion and Ka remain constant.

Since 3M is higher than 1M, the former should be regarded as the stronger acid.

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