Organic Chemistry 3

Acids and Bases for Organic Chemistry

A review from general chemistry

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Strong Acids and Bases

As a review from the first semester of general chemistry, the list of strong acids and bases are the most important as these are among the most common acids and bases. Know their formulas. Remember that the word "strong" means complete ionization.

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Acid Base Definitions

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Bronsted Lowry and Arrhenius examples

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Lewis Acid-Base example

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Acid and Base Strength

  • Strong acids are completely dissociated in water.

    • Their conjugate bases are quite weak.

  • Weak acids only dissociate partially in water.

    • Their conjugate bases are weak bases.

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Acid strength and Equilibria

  • In any acid–base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.

  • Acetate is a stronger base than H2O, so the equilibrium favors the left side (K < 1).

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The greater the value of Ka, the stronger the acid.

Organic chemists and biochemists use pKa instead of Ka

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Factors affecting acid strength

Five factors that determine acid strength (Bronsted-Lowry) will help you to understand the trend seen in the diagram on the right and above.


1 Periodicity within a column of the periodic table (equivalent to Bond Strength Effects)


2 Hybridization


3 Resonance Effects


4 Inductive Effects and


5 Electronegativity effects.


Before you get all atwitter over having to understand 5 effects, there really are only two major factors: bond strength and anion stability. With the exception we will discuss shortly, anion stability trumps bond strength. This is all related to the equilibrium of acid dissociation. Except for our six strong acids, the rest are weak which means a Ka (pKa) value. Bond strength is a reactant factor-the weaker the bond, the more the acid dissociates (higher Ka, lower pH). Anion stability is about the anion product of acid dissociation. If the anion is not stable, it will recombine with H+ and the reverse reaction dominates so there is less dissociation (lower Ka, higher pH). The more stable the anion, the more dissociation "remains" (less reverse reaction and so on).


Bond strength is about how strong the bond is to hydrogen. In order for acids to do what acids do (Arrhenius and Bronsted-Lowry), the acid must lose a hydrogen. In order for a base to come along and take the hydrogen, the bond must be weaker rather than stronger. So we see this effect most often in a column of the periodic table-as in the binary acids for the halogens. In that column, HF is not a strong acid (even though the HF bond is the most polar) but the rest are (HCl, HBr, HI) with HI being the strongest because the H-I bond is the weakest.


Anion stability is affected by the other four factors (1) resonance stabilization, (2) inductive effects, (3) hybridization effects and (4) electronegativity effects.

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Electronegativity on the central atom of the anion helps stabilize the anion

Electronegativity affects anion stability because the higher the electronegativity of the atom upon which the negative charge rests (anion), i.e., the ability of the atom to stabilize additional electrons, the more stable the anion. This effect helps explain the left to right trend for increasing acidity on the periodic table.
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Resonance helps stabilize the anion by delocalizing the negative charge

When the central atom is the same such that electronegativity does not play a role, AND when the hybridization is not that different, a series of anions can be differentially stabilized by resonance. When charge is delocalized over more than one atom, resonance occurs. An excellent example is the comparison of alcohols and carboxylic acids:



  • The conjugate base of an alcohol is an alkoxide anion (RO-). In the case of ethanol, it is ethoxide anion. In such an alkoxide ion,the negative charge is essentially localized upon oxygen. This being an electronegative atom, the negative charge is fairly stable there, so that alcohols are modestly acidic, comparable to water. This is reasonable because in the conjugate base of water, the negative charge is also located upon an oxygen atom.




  • The conjugate base of a carboxylic acid is a carboxylate anion. In the case of ,say, acetic acid, it is acetate anion. In such a carboxylate anion, the negative charge is also upon oxygen, but because of resonance effects it is delocalized over two oxygen atoms. Therefore additional stabilization of the anion results, so that carboxylic cacids are typically about 11 powers of ten more acidic than alcohols.

Hybridization exerts an even greater influence on anion stability than electronegativity.

An especially interesting case to consider is the comparison between anions in which the negative charge is on the same atom but the hybridization state of the atom is varied. The most important case here is that of carbon, which, as you know, has three distinct hybridization states.


  • The acidity of ethyne(acetylene) is much greater than than of ethene, which is much greater than than of ethane. The relevant pKa's are 25, 44, and 51. The details lie within the energies of the hybridized orbitals. Anions have an unshared pair of electrons. The energy of the orbital in which this pair of electrons resides determines the relative stability of the anion: the lower the orbital energy, the more stable the anion. sp orbitals are lowest in energy and sp3 orbitals are the highest in energy-hence, the huge difference in pKa values quoted above.

Inductive effects help to explain why chloroacetic acid is a stronger acid than acetic acid

This effect on anion stability is a complex one which involves attraction and repulsions within one molecule due to the presence of an electronegative atom like chlorine creating a dipole that is attracted (due to orientation) to the negative anion. This is a particularly strong effect as chloroacetic acid is almost 100 times more acidic than acetic acid. The closer the electronegative atom is to the carboxylate group, the more effective it is in "pulling" electron density away from the carboxylate ion, making the acid more acidic.


Alkyl groups attached to carboxylic acids have the opposite effect: these groups "push" electrons towards the carboxylate ion (remember we are discussing the stabilization of the anion) so that acetic acid (ethanoic acid) is less acidic than methanoic acid and the same effect is seen for propanoic acid and for butanoic acid and so on. pKa values rise as the number of carbons increase in the carboxylic acid.



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Oxyacid acidity is determined by the number of oxygen atoms

As you may have learned in general chemistry, in an oxyacid series, the greater the number of oxygen atoms, the stronger the acid. Let's see if we can explain this in terms of what we have learned above. The central atom is the same as in the example above for the chlorine oxyacids so electronegativity is relatively constant. The hybridization on the chlorine and oxygen atoms in this series is also relatively the same. The reasons for this trend are probably a mixture of resonance (in the anions for ClO2-, ClO3-, and ClO4- and inductive effects with the highly electronegative oxygen atoms creating dipoles with the chlorine atom.

Polyprotic acids have more than one Ka!

Polyprotic acids have more than one acidic proton.

If the difference between the Ka for the first dissociation and subsequent Ka values is 103 or more, the pH generally depends only on the first dissociation.


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Factors affecting base strength

Two of the factors which influence the strength of a base are:

  • the ease with which the lone pair picks up a hydrogen ion,

  • the stability of the ions being formed.


This list should look familiar as it essentially mirrors the two main factors that affect acid strength. The difference is that instead of bond strength, the first factor listed above is about how negative the central atom is to begin with so that a H+ is more highly attracted. SO, the inductive effect of alkyl groups pushing electron density toward the central atom actually helps methylammonia be more basic than ammonia. AND the cation formed, methylammonium, is able to spread charge around more effectively due to the electron pushing alkyl group making the cation formed more stable. The strength of weak bases, like weak acids, mainly depends on ion stability (anion or cation) to push the weak equilibrium more to the right increasing either acidity or basicity.


In acids, resonance stabilizes the anion making the acid more acidic. In bases, like aniline, this effect works to make aromatic bases like aniline less basic. The delocalization of the benzene ring is too close to the nitrogen atom in aniline. Therefore, the lone pair on nitrogen gets caught up in the delocalization to some extent and LOWERS the ability of the nitrogen to attract a H+ ion.

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