Like water, alcohols are weak acids and bases. This is attributed to the polarization of the O–H bond making the hydrogen partially positive. Moreover, the electron pairs on the oxygen atom of alcohol make it both basic and nucleophilic. Protonation of an alcohol converts hydroxide, a poor leaving group, into water—a good one. The two acid–base equilibria corresponding to ethanol are depicted below.

Figure 1. Loss of proton

Figure 2. Gain of proton

Methanol (pK_a = 15.5) is the only alcohol that is slightly stronger than water (pK_a = 15.7). Ethanol (pK_a = 15.9), tert-butanol (pK_a = 18.0), and others are weaker acids. However, all alcohols are stronger acids than terminal alkynes, and they are much stronger than hydrogen, ammonia, and alkanes.

Figure 3. Relative acidity

Since alcohols are weaker acids than water, their conjugate base alkoxide ions are stronger bases than the hydroxide ion. Alcohol can be converted into metal alkoxide using strong bases such as sodium/potassium hydride or sodium/potassium metal, which react violently but controllably with alcohol. When the alkyl substitution is bulky, the alkoxide ion is not solvated enough due to the steric effect, leading to less stabilization. Additionally, destabilization is favored by inductive effects. Consequently, the equilibrium lies predominantly towards alcohol.

Alcohols can also act as a base and accept protons from strong acids. Notably, conjugate bases of compounds with higher pK_a than an alcohol will deprotonate that alcohol.

Figure 4. Relative basicity

Phenols are more acidic than alcohols. The conjugate base of a phenol is a phenoxide or phenolate ion. Resonance stabilization of the phenoxide ion coupled with the polar effect of the benzene ring enhances the acidity of phenols by eight orders of magnitude (100,000,000 times) over cyclohexanol. Therefore, phenol does not need to be deprotonated with a base as strong as sodium hydride. Instead, it can be deprotonated by hydroxide, unlike an alcohol.

Although phenol is a million-fold higher in acidity than ethanol, it is a hundred thousand-fold less acidic than acetic acid. Their relative acid–base properties can be used to separate each other from a mixture. When an ether solution of a mixture of alcohol and phenol is extracted with dilute sodium hydroxide, phenol gets completely partitioned into the aqueous phase as its sodium salt, while alcohol stays in the ether layer. On the other hand, dilute sodium bicarbonate is used to extract phenol and carboxylic acid from an ether solution of their mixture. Carboxylic acid gets quantitatively converted into its sodium salt and gets extracted from water while phenol remains in the ether phase.

Figure 5. Acid–base equilibria of phenol

A phenoxide ion can be stabilized by delocalization of the oxygen's negative charge on the benzene ring. This is reinforced by electron-withdrawing functional groups like nitro, halide, etc. A substantial change in acidity is noted in phenols with an electron-withdrawing substituent, like a nitro group. An ortho- or para-nitro group stabilizes the phenoxide ion by delocalizing the negative charge on its own oxygen atoms. On the other hand, a meta-nitro group, being not directly conjugated to the phenoxide oxygen, stabilizes the phenolate ion to a lesser extent. Therefore, m-nitrophenol (pK_a = 8.4) is more acidic than phenol but less acidic than o- or p-nitrophenol (pK_a = 7.2). This also explains the extremely high acidity of 2,4-⁠dinitrophenol(pK_a = 4.0) and 2,4,6-trinitrophenol (pK_a = 0.4).

Table 1. The acidity constants (pK_a) of acids (blue) and their conjugate bases.

Compound	Acid	Conjugate base	pKa
Hydrogen chloride			−6.30
Nitrophenol			7.07
Phenol			9.89
m-Cresol			10.1
2,2,2-Trifluoroethanol			12.0
Water			15.7
Ethanol			15.9
Cyclohexanol			16.0