Chapter 15: Acid-Based Equilibria
Learning Outcomes
- Identify acids, bases, and conjugate acid-base pairs according to the Brønsted-Lowry definition
- Write equations for acid and base ionization reactions
- Use the ion-product constant for water to calculate hydronium and hydroxide ion concentrations
- Describe the acid-base behavior of amphiprotic substances
The acid-base reaction class has been studied for quite some time. In 1680, Robert Boyle reported traits of acid solutions that included their ability to dissolve many substances, to change the colors of certain natural dyes, and to lose these traits after coming in contact with alkali (base) solutions. In the eighteenth century, it was recognized that acids have a sour taste, react with limestone to liberate a gaseous substance (now known to be [latex]\ce{CO2}[/latex]), and interact with alkalis to form neutral substances. In 1815, Humphry Davy contributed greatly to the development of the modern acid-base concept by demonstrating that hydrogen is the essential constituent of acids. Around that same time, Joseph Louis Gay-Lussac concluded that acids are substances that can neutralize bases and that these two classes of substances can be defined only in terms of each other. The significance of hydrogen was reemphasized in 1884 when Svante Arrhenius defined an acid as a compound that dissolves in water to yield hydrogen cations (now recognized to be hydronium ions) and a base as a compound that dissolves in water to yield hydroxide anions.
Johannes Brønsted and Thomas Lowry proposed a more general description in 1923 in which acids and bases were defined in terms of the transfer of hydrogen ions, H+. (Note that these hydrogen ions are often referred to simply as protons, since that subatomic particle is the only component of cations derived from the most abundant hydrogen isotope, [latex]\ce{^1H}[/latex].) A compound that donates a proton to another compound is called a Brønsted-Lowry acid, and a compound that accepts a proton is called a Brønsted-Lowry base. An acid-base reaction is, thus, the transfer of a proton from a donor (acid) to an acceptor (base).
The concept of conjugate pairs is useful in describing Brønsted-Lowry acid-base reactions (and other reversible reactions, as well). When an acid donates [latex]\ce{H+}[/latex], the species that remains is called the conjugate base of the acid because it reacts as a proton acceptor in the reverse reaction. Likewise, when a base accepts [latex]\ce{H+}[/latex], it is converted to its conjugate acid. The reaction between water and ammonia illustrates this idea. In the forward direction, water acts as an acid by donating a proton to ammonia and subsequently becoming a hydroxide ion, [latex]\ce{OH-}[/latex], the conjugate base of water. The ammonia acts as a base in accepting this proton, becoming an ammonium ion, [latex]\ce{NH4+,NH4+}[/latex],the conjugate acid of ammonia. In the reverse direction, a hydroxide ion acts as a base in accepting a proton from ammonium ion, which acts as an acid.
The reaction between a Brønsted-Lowry acid and water is called acid ionization. For example, when hydrogen fluoride dissolves in water and ionizes, protons are transferred from hydrogen fluoride molecules to water molecules, yielding hydronium ions and fluoride ions:
Base ionization of a species occurs when it accepts protons from water molecules. In the example below, pyridine molecules, [latex]\ce{C5NH5}[/latex], undergo base ionization when dissolved in water, yielding hydroxide and pyridinium ions:
The preceding ionization reactions suggest that water may function as both a base (as in its reaction with hydrogen fluoride) and an acid (as in its reaction with ammonia). Species capable of either donating or accepting protons are called amphiprotric, or more generally, amphoteric, a term that may be used for acids and bases per definitions other than the Brønsted-Lowry one. The equations below show the two possible acid-base reactions for two amphiprotic species, bicarbonate ion and water:
[latex]\ce{HCO3-}(aq) + \ce{H2O}(l) \qquad \ce{CO3^2-} (aq) + \ce{H3O+}(aq)[/latex]
[latex]\ce{HCO3-}(aq) + \ce{H2O}(l) \qquad \ce{H2CO3}(aq) + \ce{OH-}(aq)[/latex]
The first equation represents the reaction of bicarbonate as an acid with water as a base, whereas the second represents reaction of bicarbonate as a base with water as an acid. When bicarbonate is added to water, both these equilibria are established simultaneously and the composition of the resulting solution may be determined through appropriate equilibrium calculations, as described later in this chapter.
In the liquid state, molecules of an amphiprotic substance can react with one another as illustrated for water in the equations below:
The process in which like molecules react to yield ions is called autoionization. Liquid water undergoes autoionization to a very slight extent; at 25 °C, approximately two out of every billion water molecules are ionized. The extent of the water autoionization process is reflected in the value of its equilibrium constant, the ion-product constant for water, Kw:
[latex]\ce{H2O}(l)+\ce{H2O}(l)\rightleftharpoons \ce{H3O+}(aq) + \ce{OH-}(aq)\qquad{K}_{\text{w}} = [\ce{H3O+}][\ce{OH-}][/latex]
The slight ionization of pure water is reflected in the small value of the equilibrium constant; at 25 °C, Kw has a value of 1.0 [latex]\times[/latex] 10−14. The process is endothermic, and so the extent of ionization and the resulting concentrations of hydronium ion and hydroxide ion increase with temperature. For example, at 100 °C, the value for Kw is about 5.1 [latex]\times[/latex] 10−13, roughly 100-times larger than the value at 25 °C.
Example 15.1.1: Ion Concentrations in Pure Water
What are the hydronium ion concentration and the hydroxide ion concentration in pure water at 25 °C?
[reveal-answer q=”160586″]Show Solution[/reveal-answer]
[hidden-answer a=”160586″]
The autoionization of water yields the same number of hydronium and hydroxide ions. Therefore, in pure water, [latex]\left[{\ce{H}}_{3}{\ce{O}}^{\text{+}}\right]=\left[\ce{OH}^{-}\right][/latex]. At 25 °C: [latex]{K}_{\text{w}}=\left[{\ce{H}}_{3}{\ce{O}}^{\text{+}}\right]\left[{\ce{OH}}^{-}\right]={\left[{\ce{H}}_{3}{\ce{O}}^{\text{+}}\right]}^{\text{2+}}={\left[{\ce{OH}}^{-}\right]}^{\text{2+}}=1.0\times {10}^{-14}[/latex]
So:
[latex]\left[{\ce{H}}_{3}{\ce{O}}^{\text{+}}\right]=\left[{\ce{OH}}^{-}\right]=\sqrt{1.0\times {10}^{-14}}=1.0\times {10}^{-7}M[/latex]
The hydronium ion concentration and the hydroxide ion concentration are the same, and we find that both equal 1.0 × 10−7M.
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Check Your Learning
Example 15.1.2: The Inverse Proportionality of [H3O+] and [OH−]
A solution of carbon dioxide in water has a hydronium ion concentration of 2.0 [latex]\times[/latex] 10−6M. What is the concentration of hydroxide ion at 25 °C?
[reveal-answer q=”647487″]Show Solution[/reveal-answer]
[hidden-answer a=”647487″]
We know the value of the ion-product constant for water at 25 °C:
[latex]\ce{2H2O}(l)\rightleftharpoons \ce{H3O+}(aq)+\ce{OH-}(aq){K}_{\text{w}} = [\ce{H3O+}][\ce{OH-}]=1.0\times {10}^{-14}[/latex]
Thus, we can calculate the missing equilibrium concentration.
Rearrangement of the Kw expression yields that [latex]\ce{OH-}[/latex] is directly proportional to the inverse of [latex][\ce{H3O+}][/latex]:
[latex][\ce{OH-}]=\dfrac{{K}_{\text{w}}}{[\ce{H3O+}]}=\dfrac{1.0\times {10}^{-14}}{2.0\times {10}^{-6}}=5.0\times {10}^{-9}[/latex]
The hydroxide ion concentration in water is reduced to 5.0 [latex]\times[/latex] 10−9M as the hydrogen ion concentration increases to 2.0 [latex]\times[/latex] 10−6M. This is expected from Le Châtelier’s principle; the autoionization reaction shifts to the left to reduce the stress of the increased hydronium ion concentration and the [latex]\ce{OH-}[/latex] is reduced relative to that in pure water.
A check of these concentrations confirms that our arithmetic is correct:
[latex]{K}_{\text{w}} = [\ce{H3O+}][\ce{OH-}]=\left(2.0\times {10}^{-6}\right)\left(5.0\times {10}^{-9}\right)=1.0\times {10}^{-14}[/latex]
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Check Your Learning
Example 15.1.3: Representing the Acid-Base Behavior of an Amphoteric Substance
Write separate equations representing the reaction of [latex]\ce{HSO3-}[/latex]
- as an acid with [latex]\ce{OH-}[/latex]
- as a base with [latex]\ce{HI}[/latex]
[reveal-answer q=”236960″]Show Solution[/reveal-answer]
[hidden-answer a=”236960″]
- [latex]\ce{HSO3-}(aq)+\ce{OH-}(aq)\rightleftharpoons \ce{SO3^2-}(aq)+\ce{H2O}(l)[/latex]
- [latex]\ce{HSO3-}(aq)+\ce{HI}(aq)\rightleftharpoons \ce{H2SO3}(aq)+\ce{I-}(aq)[/latex]
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Check Your Learning
Key Concepts and Summary
A compound that can donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Thus, an acid-base reaction occurs when a proton is transferred from an acid to a base, with formation of the conjugate base of the reactant acid and formation of the conjugate acid of the reactant base. Amphiprotic species can act as both proton donors and proton acceptors. Water is the most important amphiprotic species. It can form both the hydronium ion, [latex]\ce{H3O+}[/latex], and the hydroxide ion, [latex]\ce{OH-}[/latex] when it undergoes autoionization:
[latex]\ce{2H2O}(l)\rightleftharpoons \ce{H3O+}(aq)+\ce{OH-}(aq)[/latex]
The ion product of water, Kw is the equilibrium constant for the autoionization reaction:
[latex]{K}_{\text{w}}=[\ce{H2O+}][\ce{OH-}]=1.0\times 1{0}^{-14}\text{ at }25^{\circ}\text{C}[/latex]
Key Equations
- [latex]{K}_{\text{w}}=[\ce{H2O+}][\ce{OH-}]=1.0\times 1{0}^{-14}\text{ at }25^{\circ}\text{C}[/latex]
Try It
- Write equations that show [latex]\ce{NH3}[/latex] as both a conjugate acid and a conjugate base.
- Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:
- [latex]\ce{H3O+}[/latex]
- [latex]\ce{HCl}[/latex]
- [latex]\ce{NH3}[/latex]
- [latex]\ce{CH3CO2H}[/latex]
- [latex]\ce{NH4+}[/latex]
- [latex]\ce{HSO4-}[/latex]
- Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base:
- [latex]\ce{H2O}[/latex]
- [latex]\ce{OH-}[/latex]
- [latex]\ce{NH3}[/latex]
- [latex]\ce{CN-}[/latex]
- [latex]\ce{S^2-}[/latex]
- [latex]\ce{H2PO4-}[/latex]
- Identify and label the Brønsted-Lowry acid, its conjugate base, the Brønsted-Lowry base, and its conjugate acid in each of the following equations:
- [latex]\ce{HNO3}+\ce{H2O}\longrightarrow \ce{H3O+}+\ce{NO3-}[/latex]
- [latex]\ce{CN-}+\ce{H2O}\longrightarrow \ce{HCN}+\ce{OH-}[/latex]
- [latex]\ce{H2SO4}+\ce{Cl-}\longrightarrow \ce{HCl}+\ce{HSO4-}[/latex]
- [latex]\ce{HSO4-}+\ce{OH-}\longrightarrow \ce{SO4^2-}+\ce{H2O}[/latex]
- [latex]\ce{O^2-}+\ce{H2O}\longrightarrow \mathrm{2OH-}[/latex]
- [latex]{\left[\ce{Cu}{\left({\ce{H}}_{2}\ce{O}\right)}_{3}\left(\ce{OH}\right)\right]}^{\text{+}}+{\left[\text{Al}{\left({\ce{H}}_{2}\ce{O}\right)}_{6}\right]}^{3+}\longrightarrow {\left[\ce{Cu}{\left({\ce{H}}_{2}\ce{O}\right)}_{4}\right]}^{2+}+{\left[\ce{Al}{\left({\ce{H}}_{2}\ce{O}\right)}_{5}\left(\ce{OH}\right)\right]}^{2+}[/latex]
- [latex]{\ce{H}}_{2}\ce{S}+{\ce{NH}}_{2}{}^{-}\longrightarrow {\ce{HS}}^{-}+{\ce{NH}}_{3}[/latex]
[reveal-answer q=”41360″]Show Selected Solutions[/reveal-answer]
[hidden-answer a=”41360″]
1. One example for [latex]\ce{NH3}[/latex] as a conjugate acid: [latex]\ce{NH2-}+\ce{H+}\longrightarrow \ce{NH3}[/latex]; as a conjugate base: [latex]\ce{NH4+}(aq)+\ce{OH-}(aq)\longrightarrow \ce{NH3}(aq)+\ce{H2O}(l)[/latex]
2. The net ionic equations are as follows:
- [latex]\ce{H3O+}(aq)\longrightarrow \ce{H+}(aq)+\ce{H2O}(l)[/latex]
- [latex]\ce{HCl}(l)\longrightarrow \ce{H+}(aq)+\ce{Cl-}(aq)[/latex]
- [latex]\ce{NH3}(aq)\longrightarrow \ce{H+}(aq)+\ce{NH2-}(aq)[/latex]
- [latex]\ce{CH3CO2H}(aq)\longrightarrow \ce{H+}(aq)+\ce{CH3CO2-}(aq)[/latex]
- [latex]\ce{NH4+}(aq)\longrightarrow \ce{H+}(aq)+\ce{NH3}(aq)[/latex]
- [latex]\ce{HSO4-}(aq)\longrightarrow \ce{H+}(aq)+\ce{SO4^2-}(aq)[/latex]
3. The net ionic equations are as follows:
- [latex]\ce{H2O}(l)+\ce{H+}(aq)\longrightarrow \ce{H3O+}(aq)[/latex]
- [latex]\ce{OH-}(aq)+\ce{H+}(aq)\longrightarrow \ce{H2O}(l)[/latex]
- [latex]\ce{NH3}(aq)+\ce{H+}(aq)\longrightarrow \ce{NH4+}(aq)[/latex]
- [latex]\ce{CN-}(aq)+\ce{H+}(aq)\longrightarrow\ce{HCN}(aq)[/latex]
- (e) [latex]\ce{S^2-}(aq)+\ce{H+}(aq)\longrightarrow \ce{HS-}(aq)[/latex]
- (f) [latex]\ce{H2PO4-}(aq)+\ce{H+}(aq)\longrightarrow \ce{H3PO4}(aq)[/latex]
4. The labels are Brønsted-Lowry acid = BA; its conjugate base = CB; Brønsted-Lowry base = BB; its conjugate acid = CA.
- [latex]\ce{HNO3(BA)}[/latex], [latex]\ce{H2O(BB)}[/latex], [latex]\ce{H3O+}\left(\text{CA}\right)[/latex], [latex]{\text{NO}}_{3}{}^{-}\left(\text{CB}\right)[/latex]
- [latex]\ce{CN-(BB)}[/latex], [latex]\ce{H2O(BA)}[/latex], [latex]\ce{HCN(CA)}[/latex], [latex]\ce{OH-(CB)}[/latex]
- [latex]\ce{H2SO4(BA)}[/latex], [latex]\ce{Cl-(BB)}[/latex], [latex]\ce{HCl(CA)}[/latex], [latex]{\ce{HSO}}_{4}{}^{-}\left(\text{CB}\right)[/latex]
- [latex]{\ce{HSO}}_{4}{}^{-}\left(\text{BA}\right)[/latex], OH−(BB), [latex]{\ce{SO}}_{4}{}^{\text{2-}}[/latex] (CB), H2O(CA)
- [latex]\ce{O^2-(BB)}[/latex], [latex]\ce{H2O(BA)}[/latex] [latex]\ce{OH-} \text{(CB and CA)}[/latex]
- [latex]{\left[{\ce{Cu(H}}_{2}{\ce{O)}}_{3}\ce{(OH)}\right]}^{\text{+}}\text{(BB)}[/latex], [latex]{\left[{\ce{Al(H}}_{2}{\ce{O)}}_{6}\right]}^{3+}\left(\text{BA}\right)[/latex], [latex]{\left[\ce{Cu}{\left({\ce{H}}_{2}\ce{O}\right)}_{4}\right]}^{2+}\left(\text{CA}\right)[/latex], [latex]{\left[\ce{Al}{\left({\ce{H}}_{2}\ce{O}\right)}_{5}\left(\ce{OH}\right)\right]}^{2+}\left(\text{CB}\right)[/latex]
- H2S(BA), [latex]{\ce{NH}}_{2}{}^{-}\left(\text{BB}\right)[/latex], HS−(CB), NH3(CA)
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Glossary
acid ionization: reaction involving the transfer of a proton from an acid to water, yielding hydronium ions and the conjugate base of the acid
amphiprotic: species that may either gain or lose a proton in a reaction
amphoteric: species that can act as either an acid or a base
autoionization: reaction between identical species yielding ionic products; for water, this reaction involves transfer of protons to yield hydronium and hydroxide ions
base ionization: reaction involving the transfer of a proton from water to a base, yielding hydroxide ions and the conjugate acid of the base
Brønsted-Lowry acid: proton donor
Brønsted-Lowry base: proton acceptor
conjugate acid: substance formed when a base gains a proton
conjugate base: substance formed when an acid loses a proton
ion-product constant for water (Kw): equilibrium constant for the autoionization of water
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proton donor
proton acceptor
substance formed when an acid loses a proton
substance formed when a base gains a proton
reaction involving the transfer of a proton from an acid to water, yielding hydronium ions and the conjugate base of the acid
reaction involving the transfer of a proton from water to a base, yielding hydroxide ions and the conjugate acid of the base
species that may either gain or lose a proton in a reaction
species that can act as either an acid or a base
reaction between identical species yielding ionic products; for water, this reaction involves transfer of protons to yield hydronium and hydroxide ions
equilibrium constant for the autoionization of water