Drawing the conjugate bases, Oxygen is more electronegative than Chlorine so negative charge on Oxygen is more stable than Chlorine, Thus My argument that Acetic Acid is more acidic.

Also the O¯ ion is in resonance with another Oxygen so that should also add to the acidic nature.

Where am I going wrong?

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    $\begingroup$ I can't dig into my textbook right now, but I believe the reasoning is that water solvates chloride ions extremely well - so a massive amount can dissolve very easily. Acetate ions are stable, but don't interact as well with water. I'll try and get a proper answer written later this evening, if I haven't forgotten (: $\endgroup$ – etherealflux Mar 14 '16 at 15:48
  • $\begingroup$ Also acetic acid tends to dimerise, thus reducing the effective number of acetate ions $\endgroup$ – user10153 Mar 14 '16 at 17:03
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    $\begingroup$ The chloride anion and the acetate anion are such different species that the "why" in the question is hard to explain without blowing smoke. The place to start is to recognize that the Pauling electronegativity scale is for free atoms in space, not ions in aqueous solution. // Obviously in aqueous solutions you can lookup pKa's to order acid strengths. $\endgroup$ – MaxW Mar 14 '16 at 22:59

A simpler question, would be why is water a weaker acid than $\ce{HCl}$? If we were to look at that issue, we might clear some of the mess up.

Strongly reactive compounds produce weakly reactive products.

In the case of acid base theory, chloride ion is a much weaker base than hydroxide. Why is this so? The hydroxide is much smaller, and what is referred to as 'harder'. The guts of the issue is that in order for a base to deprotonate anything it has to donate two electrons to a bonding region near a proton (and in doing such render the orbital manifest, and the bond existent). In oxygen based anions, those two electrons are more likely to be in a confined region, and hence more likely to donate to the putative $\ce{HO-H}$ bond that is forming, while $\ce{Cl^-}$, being larger has a harder time getting two electrons into the bonding space required to form a $\ce{Cl-H}$ bond. So hydroxide is a stronger base. This means that when donating a proton, it is likely to just yank it back again from whatever it just gave it up to, as compared to $\ce{Cl^-}$. $\ce{HCl}$ gives up the proton and is more likely to not recapture it.

Groovy? Back to acetate. The resonance in acetate creates some diffusion of the negative charge and makes it harder to form said $\ce{AcO-H}$ bond, and as such it is a weaker base than hydroxide. However, the smaller oxygen atoms are still better and forming bonds than $\ce{Cl^-}$ and as such it is a stronger base than chloride, and therefore its conjugate acid is less acidic.

Strong acids: weak conjugate bases. Strong bases: weak conjugate acids.

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    $\begingroup$ Yep. Chloride ions are incredibly terrible bases, which makes their conjugate acid (HCl) very powerful. $\endgroup$ – etherealflux Mar 15 '16 at 0:36
  • $\begingroup$ "Strong acids: weak conjugate bases. Strong bases: weak conjugate acids." This is really more of fact that could be derived from pKa values than an "explanation." In aqueous solution: $$\mathrm{pKa + pKb = 14}$$ $\endgroup$ – MaxW Mar 15 '16 at 22:01

Acid strength depends upon ease of ionization of the 'acidic hydrogen' forming hydronium ions by proton transfer. (Bronsted-Lowry Theory). HCl is one of the 'strong acids' and as such ionizes 100% whereas Acetic Acid (HOAc <=> H + OAc) is a 'weak acid' and ionizes less than 100% (~1.3% at 25C). For a 0.10M HCl(aq) solution => 0.10M in hydronium ions, but a 0.10M HOAc(aq) solution => only 0.0013M in hydronium ions at 25C. Structurally, the electronegativity of the chloride ion is much less than that of the acetate ion which means the covalent bond of the ionizable hydrogen in HCl is much weaker and more easily ionized in aqueous solution than the ionizable hydrogen to oxygen bond in HOAc.

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Acid strength is determined by the stability of the conjugate base. Here are the factors that influence stability of the conjugate base:

  1. Ion size (when going down the rows of the periodic table)
  2. Electronegativity (when going in the same row of the periodic table across)
  3. Resonance
  4. Inductive Effects

If you look at the conjugate bases, in acetic acid, the negative charge falls on the oxygen and in HCl, the negative charge falls on the chlorine. Because of the fact that oxygen and chlorine are on different rows, we use ion size as the most significant factor for determining CB stability. The size of the chloride ion is greater than the size of the oxygen ion and therefore the CB of HCl can distribute the charge more evenly and is thus more stable. Moreover, the methyl group in the acetic acid is also electron-donating, meaning that it actually reduces how much the charge gets spread around on the molecule’s conjugate base. However, this is not really a significant factor.

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