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Comparison of the strengths of weak acids
The strength of weak acids is measured by pKaAScale. The lower the number on this scale, the stronger the acid.
Let us consider three of the compounds together with their pKaAThe values are:
Remember: the lower the number, the stronger the acid. If you compare the other two with acetic acid you will see that phenol is much weaker with a pKaAthan 10.00, and ethanol is as weak with a pKaAof around 16 which is hardly considered acid!
Acidity of carboxylic acids
about pKASome typical carboxylic acids are listed in the table below. Comparing these values with those of comparable alcohols such as ethanol (pKA= 16) e 2-Metil-2-propanol (pKA= 19), it is clear that carboxylic acids are more than ten powers of ten strongest acids! Additionally, electronegative substituents near the carboxyl group act to increase acidity.
| Connection | packageA | Connection | packageA | |
|---|---|---|---|---|
| HCO2H | 3,75 | CH3CH2CH2CO2H | 4,82 | |
| CH3CO2H | 4,74 | ClCH2CH2CH2CO2H | 4.53 | |
| FCH2CO2H | 2,65 | CH3CHClCH2CO2H | 4.05 | |
| ClCH2CO2H | 2,85 | CH3CH2CHClCO2H | 2,89 | |
| BrCH2CO2H | 2,90 | C6H5CO2H | 4.20 | |
| UE2CO2H | 3.10 | post office2NORTH CAROLINA6H4CO2H | 3,45 | |
| Kl3CCO2H | 0,77 | p-CH3OK6H4CO2H | 4,45 |
Why would the presence of a carbonyl group next to a hydroxyl group have such a profound effect on the acidity of the hydroxyl proton? To answer this question, we must go back to the nature of acid-base balance and the definition of pKa.A, illustrated by the general equations given below. These relationships were described in aprevious sectionthat text
We know that an equilibrium favors the thermodynamically more stable side and that the magnitude of the equilibrium constant reflects the difference in energy between the components on each side. In an acid-base balance, the balance always favors the weaker acid and base (which are the more stable components). Water is the default base used for pKAmeasurements; therefore, anything that stabilizes the conjugate base (A:(-)) of an acid, this acid (H-A) inevitably becomes stronger and shifts the equilibrium to the right. Both the carboxyl group and the carboxylate anion are resonance stabilized, but the stabilization of the anion is much greater than that of the neutral function, as shown in the diagram below. For the carboxylate anion, the two contributing structures in the hybrid have the same weight and the C-O bonds are the same length (between a double bond and a single bond). This stabilization leads to a significant increase in acidity, as illustrated in the energy diagram which can be accessed by clicking on the button "toggle screen" I like it.
Compounds such as alcohols and phenol that contain an -OH group attached to a hydrocarbon are very weak acids. Alcohols are so weakly acidic that, for normal laboratory purposes, their acidity is practically negligible. However, phenol is acidic enough to have recognizable acidic properties, although it is still a very weak acid. A hydrogen ion can be separated from the -OH group and transferred to a base.
The pKa of ethanol is around 17, while the pKa of acetic acid is around 5: this is a value of 1012times the difference between the two acid constants. In both compounds, the acidic proton is attached to an oxygen atom. How can they be so different in terms of acidity?
We begin by considering the conjugate bases.
In both types, the negative charge of the conjugate base is held by an oxygen, so periodic polarization cannot be induced. For acetic acid, however, there is a fundamental difference: a resonance contribution can be made where the negative charge is on the second oxygen in the group. The two resonance forms of the conjugate base are energetically equal. This means that the negative charge on the acetate ion is not due to oxygen, but is shared between the two. Chemists use the term "charge displacement" to describe this phenomenon. With the ethoxide ion, on the other hand, the negative charge is "stuck" on the individual oxygen, it has nowhere to go.
Remember the basic idea that electrostatic charges, whether positive or negative, are more stable when "scattered" than when confined to an atom. Here, a charge is "distributed" (i.e. discharged)by resonance, and not simply because of the size of the atom involved.
Resonant charge delocalization has a very potent effect on the reactivity of organic molecules, enough to account for the difference of more than 12 pKAUnits between ethanol and acetic acid. The acetate ion is much more stable than the ethoxide ion, all due to resonance delocalization effects.
The resonance effect also explains why a nitrogen atom is so much more basic when it is in an amine, butNoclearly basic when it is part of an amide group. Recall that in an amide there is significant double bond character to the carbon-nitrogen bond due to a second resonance contribution where the nitrogen pair is part of a p bond.
Whereas the lone pair of electrons on an amine nitrogen is "stuck" in one place, the lone pair on an amide nitrogen is resonance delocalized. Note that in this case we have extended our core statement to the effect that electron density, in the form of a lone pair of electrons, is stabilized by resonance delocalization, even though no negative charge is involved. Here's another way to think about it: The lone pair of electrons on an amide nitrogen is not available to bond with a proton; these two electrons are very "comfortable" as they are part of the delocalized pi bonding system. By contrast, the lone pair of electrons on an amine nitrogen is not part of a delocalized pi system and is quite willing to bond with any acidic protons that may be nearby.
Why is phenol acidic?
Compounds such as alcohols and phenols that contain an -OH group attached to a hydrocarbon are very weak acids. Alcohols are so weakly acidic that, for normal laboratory purposes, their acidity is practically negligible. However, phenol is acidic enough to have recognizable acidic properties, although it is still a very weak acid. A hydrogen ion can be separated from the -OH group and transferred to a base. For example dissolved in water:
Phenol is a very weak acid, and the equilibrium position is to the far left. Phenol can lose a hydrogen ion because the phenoxide ion formed is stabilized to a certain extent. The negative charge of the oxygen atom is delocalized around the ring. The more stable the ion, the greater the probability of its formation. One of the lone pairs of electrons on the oxygen atom overlaps with the delocalized electrons on the benzene ring.
This overlap results in a delocalization that extends from the ring to the oxygen atom. As a result, the negative charge is no longer completely on the oxygen, but is distributed throughout the ion.
Distributing the charge makes the ion more stable than it would be if all the charge remained on the oxygen. However, oxygen is the most electronegative element in the ion, and delocalized electrons are attracted there. This means that there is still a lot of charge around the oxygen, which tends to attract the hydrogen ion back. So phenol is just a very weak acid.
Why is phenol a much stronger acid than cyclohexanol? To answer this question, we need to evaluate how an oxygen substituent interacts with the benzene ring. As mentioned in our earlier discussion of electrophilic aromatic substitution reactions, aoxygen substituentit increases ring reactivity and favors electrophilic attack at ortho and parasitic sites. It has been suggested that resonance delocalization of a non-oxygen bonding pair of electrons in the pi electron system of the aromatic ring is responsible for this substituent effect. A similar set of resonance structures for the conjugate base of the phenolate anion appears below the phenolic structures.
The resonance stabilization is very different in these two cases. an importantresonance principleis that charge separation reduces the importance of canonical contributions to the resonance hybrid and reduces overall stabilization. All the structures that contribute to the phenolic hybrid undergo charge separation, which results in a very modest stabilization of this compound. On the other hand, the phenolate anion is already charged and canonical contributors act to distribute the charge, resulting in a substantial stabilization of this species. The conjugate bases of simple alcohols are not stabilized by charge delocalization, so the acidity of these compounds is similar to that of water. To the right is an energy diagram showing the effect of resonance on cyclohexanol and phenolic acids. Since the resonance stabilization of the conjugated phenolate base is much greater than the stabilization of phenol itself, the acidity of phenol relative to cyclohexanol is increased. Evidence supporting that the negative charge of phenolate is delocalized to the ortho and para carbons of the benzene ring comes from the influence of electron-withdrawing substituents at these sites.
William Reusch, Professor Emeritus (michigan state u.),organic chemistry virtual book
jim clark (Chemguide.es)