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Alcohols, Epoxides and Ethers

By James Ashenhurst

Alcohols (3) – Acidity and Basicity

Last updated: May 21st, 2019

In the last post we said that one of the keys to the reactions of alcohols as we go forward is that the conjugate acid is a better leaving group and the conjugate base is a better nucleophile.

We might have gotten a little ahead of ourselves broaching that topic, because we haven’t yet really revisited some of the fundamentals about alcohol acidity and basicity.

That’s the purpose of today’s post! In this post we’re going to:

  1. Review 4 key points about acid base reactions
  2. Give 2 examples of alcohol acid-base reactions
  3. Review the key factors which determine the acidity
  4. Show examples of how this applies to alcohols.
  5. Give a practice question at the end.

Let’s go!

1. Four Key Points To Review About Acid Base Reactions

  1. Every acid-base reaction has 4 components: an acid, a base, a conjugate acid, and a conjugate base.When an acid loses a proton, it becomes its conjugate base. When a base gains a proton, it becomes its conjugate acid. As mentioned in the previous post, the conjugate bas of an alcohol is called an alkoxide. The conjugate acid of an alcohol is called an oxonium ion.
  2. We usually describe acid-base reactions as an equilibrium. In acid-base reactions, the equilibrium will favor the direction where a stronger acid and stronger base produces a weaker acid and a weaker base.When you add HCl to NaOH, a violent acid-base reaction occurs, which leads to the formation of H2O (a weaker acid than HCl) and NaCl (a weaker base than NaOH). As you’ve no doubt discovered when adding table salt (NaCl) to water, this reaction doesn’t proceed to any significant extent in the reverse direction.
  3. We measure acidity using a term called pKa. This is a measure of the equilibrium constant for a species giving up a proton to form its conjugate base.pKa is on a scale of about -10 to 50. Sixty orders of magnitude! The higher the pKa the less acidic it is.  Lower pKa (more negative ) = more acidic.
    Water (pKa of 15.7) is a weaker acid than HCl (pKa of -8).
  4.  The stronger the acid, the weaker the conjugate base. The weaker the acid, the stronger the conjugate base. The conjugate base of the strong acid HCl (pKa -8) is the innocuous chloride ion (Cl-), a very weak baseThe conjugate base of the weak acid H2O (pKa 15.7) is the strongly basic hydroxide ion (HO-).

2. Examples of Acid-Base Reactions Of Alcohols

Here’s an example of a favorable acid-base reaction of alcohols. Note how we’re going from a stronger acid and stronger base to a weaker acid and weaker base [pKa values tell us for sure] Here, deprotonation is very favourable. Note that the conjugate base of an alcohol is called an alkoxide.


Here’s an example of a (very) unfavorable acid-base reaction of alcohols: protonation of an alcohol by NH3. The most important reason why this is unfavourable is because we’re going from a weaker acid (pKa 38) and weaker base to a stronger acid (pKa -2) and stronger base. The equilibrium constant is about 40 orders of magnitude in the wrong direction!


3. The Key Factors Which Determine Acidity

What determines how acidic a molecule is, anyway?

The key factor in determining acidity is the stability of the conjugate base. Any factor which makes the conjugate base more stable will increase the acidity of the acid.

What does that mean, exactly? Usually, it means stabilizing negative charge since the conjugate base will always be one unit of charge more “negative” than the acid.

How is negative charge stabilized? Two ways.

  • First, by bringing the charge closer to the positively charged nucleus [“opposite charges attract”, remember]. Across a row of the periodic table, for example, basicity decreases as we go from H3C to H2N to HO to F  because the electronegativity of the atom is increasing. That negative charge is being held closer to the nucleus, and therefore is more stable. A good rule of thumb is, “the more stable a lone pair, the less basic it is. This is also why certain species are made acidic by adjacent electron-withdrawing groups.
  • Second, by spreading charge out over a larger volume. Diffuse charge is more stable than concentrated charge. Down a row of the periodic table, for example, basicity decreases as we go from F to Cl to Br to I– because that negative charge is being spread out over a larger volume (larger atoms). The larger atoms are said to be more “polarizable”. [Note that this effect dominates rather than electronegativity in this case.] This is also why resonance serves to stabilize charges; the charge is being spread across multiple atoms, therefore reducing individual charge density.

4. Applying These Principles To The Acidity Of Alcohols

How do these principles relate to alcohols? It’s quite simple, actually. Since we’ll always be comparing the same atom (oxygen) we don’t need to worry about periodic trends, and we just need to focus on resonance and adjacent electron-withdrawing groups.

Alcohols where the conjugate base is resonance stabilized will be more acidic. The classic example is cyclohexanol and phenol.

Cyclohexanol has the pKa of a typical alcohol (about 16). The pKa of phenol, however, is about 10. Let’s look:


See how that negative charge on the oxygen of phenol can be “delocalized” back into the ring? That means the charge can be spread out throughout the molecule, which is stabilizing. Any factor which stabilizes the conjugate base will increase acidity. 

4-phenolate resonance

Here’s another example. Compare ethanol (pKa 16) to 2,2,2-trifluoroethanol (pKa about 12). Why do you think trifluoroethanol is more acidic?

5-ethanol tfe

Compare their conjugate bases. What is fluorine doing here to make the conjugate base more stable?

This is an example of an inductive effect. Fluorine, being highly electronegative, pulls electron density away from the neighbouring carbon. That carbon, now being electron poor, pulls electron density away from the carbon next door. And that carbon, being slightly electron poor, can pull some electron density away from the oxygen.

The net result is that the oxygen has lower electron density, which is stabilizing. Again, stabilize the conjugate base –>  increase acidity. 

This also works if we compare alcohol variations where we change the distance between the OH and the fluorine atom.


That’s because the inductive effect decreases in magnitude the farther away we go from the electronegative atom.

We can also use electronegativity trends to determine the order of acidity in these molecules. Since fluorine is more electronegative than chlorine which is more electronegative than bromine which is more electronegative than iodine, the inductive effect will be highest for CF3 and lowest for CI3.

7-ewg halides

Finally, one last example. We can even think of examples where these two effects are combined:


Which do you think might be most acidic here?


Now that we’ve covered the key factors governing the acidity of alcohols, we’re more prepared to get into the nitty gritty of their different reactions. In the next post we’ll start discussing how acidity and basicity affects the reaction conditions we can use.

For alcohols, since we’re always dealing with oxygen, the only relevant factors here are resonance and electron withdrawing groups.

Questions? Post ’em in the comments!

Next Post – The Williamson Ether Synthesis 

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Comment section

17 thoughts on “Alcohols (3) – Acidity and Basicity

  1. I apologize if I’m wrong, but isn’t the pKa of ethanol closer to 16 (15.9 is the most cited value)?
    Also, “the electronegativity of the nucleus is increasing” and “the negatively charged oxygen on trifluoroethanol will have some of its charge redistributed to fluorine” might be somewhat difficult to understand to undergraduate students. Maybe a little bit of elaboration and/or rephrasing might be in order.

  2. Regarding the phenol vs cyclohexanol example, since most resonance forms break the aromaticity of the ring, is the charge delocalization still significant enough as to make it a more stable conjugate base?

  3. Amazing. Reopen my mind of organic chem.
    I have a question. Do you think the acidity/basicity is correlated to its 1H NMR chemical shift? Specifically for protonated amino acids.

  4. the order on your second to last rank-the-acidity is incorrect. it would be correct if electronegativity were the only player in ranking acidity, but it is not. the size of the halogen also factors in here…spreading out the charge increases acidity…i.e. why HF is a weak acid…<HCl< HBr < HI …consequently, your ranking of 2,2,2trifolorethanol should be reversed…2,2,2trifolorethanol (weakest) < 2,2,2tricholor etc… pka for triflorethanol is ~12.46..pka for tricholorethanol is ~12.02

    1. OK, so I am curious here.. I posted a question regarding electronegativity vs polarizability. I can understand polarizability of the larger halogens generating the stronger acid through the creation of the larger dipole moment or the fact of the small halogen and it’s higher electronegativity (closeness of opposite charges) withdrawing electrons more strongly, but what conditions (possibly such as solvent) will determine reactivity of the species in concern ?? I see it said that polarizability takes presidence yet as you have stated he shows opposing acidity trends in the example as you have shown. Are there environmental conditions I am overlooking that aren’t stated to justify this ?? Or maybe the characteristics oh the double bonded oxygen contributing ??

      1. The person you replied to is correct on one point and wrong on everything else. 2,2,2-trichloroethanol is indeed more acidic than 2,2,2-trifluoroethanol.

        However, the difference is very small: 12.24 vs 12.46 pka, respectively.

        Furthermore, aside from 2,2,2-trichloroethanol, the pattern is in fact correct. It should not be wholly reversed like the person you replied to stated. 2,2,2-tribromoethanol clocks in at 12.7 pka, in accordance with the pattern, and 2,2,2-triiodoethanol hasn’t been measured as far as I could find.

        Finally, the mechanism at play here has nothing to do with the polarizability in the context you’re used to or anything related to the halogenic acids, HF, HCl, HBr, and HI. The polarizability of larger molecules is relevant for you in explaining intermolecular forces (induced dipole attractions). The stability of halogen anions (the conjugate base of a halogenic acid) has to do with the charge density. Fluorine is very small, so carrying a negative charge by itself concentrates the negative charge in a very small space alongside other electrons, which leads to a repulsive and destabilizing interaction. In iodine, the single extra electron is spread out over a much larger volume which minimizes destabilizing interactions. This along with orbital overlap (HSAB theory – traditionally covered in your first inorganic chemistry course) should more or less account for the differences in halogenic acid pka.

        Honestly, just follow this page’s advice and ignore the technicality on 2,2,2-trichloroethanol. The difference of 0.3 pka is so trivially small that it would involve a complex analysis of a multitude of variables to pinpoint why, so there’s no way it pops up on an exam with the expectation you know about the exception of 2,2,2-trichloroethanol.

  5. Question.. besides solvent influence how exactly does electronegativity suddenly take precedence over polarizability in stabilizing negative charge and increasing acidity of the hydroxyl function ?? I am confused here. I understand polarizability as taking precedence due stability of the distributed charge of the larger atom and can further justify polarizability and it’s effect of charge stabilization here through the fact of the longer time required for electrons to travel around the larger atom creating a larger dipole moment and therefore a greater partial charge. Yet starting a new sentence, the higher electronegativity (due higher attraction generated through closness of opposing charges) of the smaller halides mainly in question here can b justified to be electron withdrawing to carbon and oxygen there fore stabilizing the anion from both points being electronegativity, and polarizability.. which i guess answers my question but raises then where does polarizability actually take precedence ??

    1. Any meaningful way to spread out the charge density (ex. resonance or “polarizability”) takes precedence.

      For example, phenol (the right-most molecule in question before the conclusion above), has a pka of 10 due to resonance structures. That’s several hundred times more acidic than 2,2,2-trifluoroethanol, the alcohol with 3 fluorines used as an example for the inductive effect above. CF3 is a lot of inductive power, but resonance is still more important.

      Another example for polarizability: primary alcohols (OH group) have a pka around 16 while primary thiols (SH group) have a pka of around 10 — a million times more acidic.

      Just be careful if your “polarizability of an electron cloud” is referring to stabilization by e.g. alkyl chains (I’m pretty sure this is actually just hyperconjugation), as opposed to the “polarizability” stabilizing an iodide or sulfide conjugate base (pretty sure this is just a shell or size effect). The former effect is weak (+-1 pka) and its trend reverses in water vs gas phase (e.g. H2O > methanol > ethanol), because solvent effects overpower it. Trends from the latter effect (e.g. halogenic acid strength or H2O vs H2S) are not reversed between water and gas phase.

  6. You say that, in acidity of alcohols; more stabilization of the alkoxide ion formed (by inductive effects, etc), means more acidic is the alcohol.
    Sure, but the trend you would get is not analogous to the data found in the gas phase oppose the data found from the solutions.

    The acidity of alcohols is mainly due to polarizibility and solvation.


    Organic Chemistry (Sixth Edition) by Robert Thornton Morrison & Robert Neilson Boyd [Pg. 228]

    1. Hi – I make the assumption that we are dealing with solution-phase chemistry, not gas-phase. Any quoted pKa values are quoted in water and/or DMSO as solvent.

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