65 Strength of Bases

Strong Bases

Strong bases either dissociate completely in solution to yield hydroxide ions, or deprotonate water to yield hydroxide ions.

As discussed in the previous concepts on bases, a base is a substance that can: donate hydroxide ions in solution (Arrhenius definition); accept [latex]\text{H}^+[/latex] ions (protons) (Bronsted-Lowry definition); or donate a pair of valence electrons (Lewis definition). In water, basic solutions have a pH higher than 7.0, indicating a greater concentration of [latex]\text{OH}^-[/latex] than [latex]\text{H}^+[/latex].

Strong Arrhenius Bases

A strong Arrhenius base, like a strong acid, is a compound that ionizes completely or near-completely in solution. Therefore, the concentration of hydroxide ions in a strongly basic solution is equal to that of the undissociated base. Common examples of strong Arrhenius bases are the hydroxides of alkali metals and alkaline earth metals such as [latex]\text{NaOH}[/latex] and [latex]\text{CA(OH)}_2[/latex]. Strong bases are capable of deprotonating weak acids; very strong bases can deprotonate very weakly acidic [latex]\text{C-H}[/latex] groups in the absence of water.

Some common strong Arrhenius bases include:

• Potassium hydroxide ([latex]\text{KOH}[/latex])
• Sodium hydroxide ([latex]\text{NaOH}[/latex])
• Barium hydroxide ([latex]\text{Ba(OH)}_2[/latex])
• Caesium hydroxide ([latex]\text{CsOH}[/latex])
• Strontium hydroxide ([latex]\text{Sr(OH)}_2[/latex])
• Calcium hydroxide ([latex]\text{Ca(OH)}_2[/latex])
• Lithium hydroxide ([latex]\text{LiOH}[/latex])
• Rubidium hydroxide ([latex]\text{RbOH}[/latex])

The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals). Generally, the alkali metal bases are stronger than the alkaline earth metal bases, which are less soluble. When writing out the dissociation equation of a strong base, assume that the reverse reaction does not occur, because the conjugate acid of a strong base is very weak.

Superbases (Lewis bases)

Group 1 salts of carbanions (such as butyllithium, [latex]\text{LiC}_4\text{H}_9[/latex], which dissociates into [latex]\text{Li}^+[/latex] and the carbanion [latex]\text{C}_4\text{H}_9^-[/latex]), amides ([latex]\text{NH}_2^-[/latex]), and hydrides ([latex]\text{H}^-[/latex]) tend to be even stronger bases due to the extreme weakness of their conjugate acids—stable hydrocarbons, amines, and hydrogen gas. Usually, these bases are created by adding pure alkali metals in their neutral state, such as sodium, to the conjugate acid. They are called superbases, because it is not possible to keep them in aqueous solution; this is due to the fact they will react completely with water, deprotonating it to the fullest extent possible. For example, the ethoxide ion (conjugate base of ethanol) will undergo this reaction in the presence of water:

[latex]\text{CH}_3\text{CH}_2\text{O}^- + \text{H}_2\text{O} \rightarrow \text{CH}_3\text{CH}_2\text{OH} + \text{OH}^-[/latex]

Unlike weak bases, which exist in equilibrium with their conjugate acids, the strong base reacts completely with water, and none of the original anion remains after the base is added to solution. Some other superbases include:

• Butyl lithium (n-[latex]\text{BuLi}[/latex])
• Lithium diisopropylamide (LDA) ([latex]\text{C}_6\text{H}_14\text{LiN}[/latex])
• Lithium diethylamide ([latex]\text{LDEA}[/latex])
• Sodium amide ([latex]\text{NaNH}_2[/latex])
• Sodium hydride ([latex]\text{NaH}[/latex])
• Lithium bis(trimethylsilyl)amide, ([latex]\text{(CH}_3\text{)}_3\text{Si)}_2\text{NLi}[/latex]

Superbases such as the ones listed above are commonly used as reagents in organic laboratories.

Weak Bases

In aqueous solution, a weak base reacts incompletely with water to yield hydroxide ions.

A base is a substance that can accept hydrogen ions ([latex]\text{H}^+[/latex]) or, more generally, donate a pair of valence electrons. A weak base is a chemical base that does not ionize fully in an aqueous solution. As Brønsted-Lowry bases are proton acceptors, a weak base may also be defined as a chemical base with incomplete protonation. A general formula for base behavior is as follows:

[latex]\text{B} (aq) + \text{H}_2\text{O} (aq) \rightleftharpoons \text{BH}^+ (aq) + \text{OH}^- (aq)[/latex]

A base can either accept protons from water molecules or donate hydroxide ions to a solution. Both actions raise the pH of the solution by decreasing the concentration of [latex]\text{H}^+[/latex] ions. This results in a relatively low pH compared to that of strong bases. The pH of bases in aqueous solution ranges from greater than 7 (the pH of pure water) to 14 (though some bases have pH values greater than 14). The formula for pH is:

[latex]\text{pH} = \text{-log}_{10} [ \text{H}^+ ][/latex]

Sometimes, however, it is more convenient to focus on the [latex]\text{pOH}[/latex] of bases, rather than the pH. The [latex]\text{pOH}[/latex] more directly references the ([latex]\text{OH}^-[/latex]].

[latex]\text{pOH} = \text{-log}_{10} [ \text{OH}^- ][/latex]

Some common weak bases and their corresponding [latex]\text{pK}_b[/latex] values include:

• [latex]\text{C}_6\text{H}_5\text{NH}_2[/latex] (9.38)
• [latex]\text{NH}_3[/latex] (4.75)
• [latex]\text{CH}_3\text{NH}_2[/latex] (3.36)
• [latex]\text{CH}_3\text{CH}_2\text{NH}_2[/latex] (3.27)

Smaller [latex]\text{pK}_b[/latex] values indicate higher values of [latex]\text{K}_b[/latex]_; this also indicates a stronger base.