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What prevents weak acids from dissociating completely in water?

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    $\begingroup$ I'm voting to close this question as off-topic because a lack of evidence of research effort. $\endgroup$ – user15489 Apr 26 '15 at 5:25
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    $\begingroup$ possible duplicate of Why are weak acids weak? $\endgroup$ – Martin - マーチン Apr 26 '15 at 7:29
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It seems that your fundamental misunderstanding is that some molecules will dissociate and some will not. Who is to say that they do not dissociate? In fact, virtually every single molecule will dissociate in solution, BUT here's the catch: the same holds for protonation of the conjugate base. In reality, these molecules are all constantly dissociating and reprotonating, and the relative stability of the acid and conjugate base, represented by equilibrium constants, determines the rate of the forward and backward reactions, which determines where the forward and backward reactions proceed at the same rate. This is the fundamental concept of equilibrium in a reversible process, that the system is constantly in flux, but there is a point where the overall concentrations remain constant.

What we term "strong acids" are simply acids with pKa values so low that they completely dissociate and any time a conjugate base is reprotonated it almost immediately dissociates again. There are also classes of acids that are "near strong acids" and "weak acids" with increasing pKa values.

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To give a straight answer: intramolecular forces.

Take two examples, $\ce{HCl}$ and $\ce{HF}$. We all know that $\ce{HCl}$ is a strong acid and it dissociates completely, while $\ce{HF}$ doesn't. This is because of the strength of Fluorine and Chlorine differs. Fluorine is strongly bonded to Hydrogen so less of it dissociates.

Another reason is position which is mostly found in oxoacids.

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  • $\begingroup$ I'll format it later, don't have time. $\endgroup$ – Asker123 Apr 26 '15 at 2:13
  • $\begingroup$ Ok so the $HF$ intramolecular forces are strong, but this doesn't explain why some $HF$ molecules dissociate and some don't. $\endgroup$ – Joshua Benabou Apr 26 '15 at 2:41
  • $\begingroup$ I answered that part for you! $\endgroup$ – Stagg C. Apr 26 '15 at 7:53
  • $\begingroup$ HF is almost completely dissociated, but fluoride binds hydronium. $\endgroup$ – Mithoron Apr 26 '15 at 18:03
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The Acid (HA), which calls as weak, will produce an ion (A-) by dissociation. According to Bronsted-Lowry acid-base theory, the produced ion (A-), which is called as Conjugate base of acid (HA), is unstable when its respective acid (HA) is weaker one. Because of the instability, the conjugate base (A-) involves in backward directive reaction to form its parental acid (HA) which is in non-dissociated form. In other words, there could be equilibrium in between non-dissociated acid (HA) (reactant in forward direction because of the ability to donate H+ ion) and conjugate base (A-) (reactant in backward direction because of the instability).

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That is the reason there is no complete dissociation occurs in weak acids. Apart from this, there could be the situation in which the weaker one is comparative term. I mean, in the presence of strong acid (H2SO4) the weak acid (CH3COOH) is in the state of non-dissociation form that is because of ‘Common-ion effect’. In this case also there is no complete dissociation from weak acid.

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according to ostwald dilution law, there is relationship between dissociation constant and degree of dissociation of weak electrolyte. Since, weak acids have very small dissociation constant ('k' value ) , so they do not donate all of its hydrogen ion (H+).That's why , they do not dissociate completely.

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