Molecules or Ions That Can Alternately Behave as Either a Brãƒâ¸nsted-lowry Acid or Base Are Called

Learning Objectives

Past the cease of this section, you will exist able to:

• Place acids, bases, and conjugate acid-base of operations pairs co-ordinate to the Brønsted-Lowry definition
• Write equations for acrid and base ionization reactions
• Apply the ion-production constant for water to summate hydronium and hydroxide ion concentrations
• Describe the acid-base of operations 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 alter the colors of certain natural dyes, and to lose these traits after coming in contact with brine (base of operations) solutions. In the eighteenth century, it was recognized that acids accept a sour taste, react with limestone to liberate a gaseous substance (at present known to be CO₂), and interact with alkalis to form neutral substances. In 1815, Humphry Davy contributed greatly to the development of the modern acid-base of operations concept by demonstrating that hydrogen is the essential constituent of acids. Around that same time, Joseph Louis Gay-Lussac ended that acids are substances that can neutralize bases and that these 2 classes of substances can be divers only in terms of each other. The significance of hydrogen was reemphasized in 1884 when Svante Arrhenius defined an acrid as a chemical compound that dissolves in water to yield hydrogen cations (now recognized to be hydronium ions) and a base equally 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⁺. (Annotation that these hydrogen ions are oftentimes referred to simply as protons, since that subatomic particle is the simply component of cations derived from the most abundant hydrogen isotope, ¹H.) A compound that donates a proton to another chemical compound is called a Brønsted-Lowry acid, and a compound that accepts a proton is called a Brønsted-Lowry base of operations. 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, likewise). When an acid donates H⁺, the species that remains is chosen the conjugate base of the acrid considering information technology reacts as a proton acceptor in the opposite reaction. Too, when a base accepts H⁺, information technology is converted to its conjugate acid. The reaction betwixt water and ammonia illustrates this idea. In the forward direction, water acts as an acid by donating a proton to ammonia and subsequently condign a hydroxide ion, OH⁻, the conjugate base of water. The ammonia acts every bit a base in accepting this proton, becoming an ammonium ion, ${{\text{NH}}_4}^{\text{+}}\text{,}$ the conjugate acid of ammonia. In the reverse direction, a hydroxide ion acts as a base of operations 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 of operations ionization of a species occurs when information technology accepts protons from water molecules. In the example below, pyridine molecules, C₅NH₅, 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 of operations (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 chosen amphiprotric, or more mostly, amphoteric, a term that may be used for acids and bases per definitions other than the Brønsted-Lowry i. The equations below show the two possible acid-base reactions for ii amphiprotic species, bicarbonate ion and water:

$${\text{HCO}}_{\mathrm{three}}{}^{–}\left(\text{aq}\right)+{\text{H}}_2\text{O}\left(\text{fifty}\right){\text{CO}}_{\mathrm{iii}}{}^{two–}\left(\text{aq}\right)+{\text{H}}_3{\text{O}}^{+}\left(\text{aq}\right)$$

$${\text{HCO}}_3{}^{–}\left(\text{aq}\right)+{\text{H}}_2\text{O}\left(\text{50}\right){\text{H}}_{\mathrm{two}}{\text{CO}}_3\left(\text{aq}\right)+{\text{OH}}^{\mathrm{–}}\left(\text{aq}\right)$$

The first equation represents the reaction of bicarbonate as an acid with h2o every bit a base, whereas the 2d 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 exist adamant through appropriate equilibrium calculations, as described after in this affiliate.

In the liquid state, molecules of an amphiprotic substance can react with 1 another as illustrated for water in the equations below:

The process in which similar molecules react to yield ions is chosen autoionization. Liquid water undergoes autoionization to a very slight extent; at 25 °C, approximately 2 out of every billion water molecules are ionized. The extent of the water autoionization procedure is reflected in the value of its equilibrium constant, the ion-production constant for water, K _w :

$${\text{H}}_{\mathrm{two}}\text{O}(\mathrm{fifty})+{\text{H}}_{\mathrm{ii}}\text{O}(l)\rightleftharpoons {\text{H}}_3{\text{O}}^{\text{+}}(aq)+{\text{OH}}^{\text{−}}(aq)K_{\text{w}}=[{\text{H}}_{\mathrm{iii}}{\text{O}}^{\text{+}}][{\text{OH}}^{\text{−}}]$$

The slight ionization of pure water is reflected in the small value of the equilibrium constant; at 25 °C, K _w has a value of 1.0 $\times$ 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 case, at 100 °C, the value for K _westward is about v.6 $\times$ x^−xiii, roughly fifty times larger than the value at 25 °C.

Example xiv.1

Ion Concentrations in Pure Water

What are the hydronium ion concentration and the hydroxide ion concentration in pure water at 25 °C?

Solution

The autoionization of water yields the same number of hydronium and hydroxide ions. Therefore, in pure h2o, [H₃O⁺] = [OH⁻] = x. At 25 °C:

$$K_{\text{w}}=[{\text{H}}_3{\text{O}}^{\text{+}}][{\text{OH}}^{\text{−}}]=\left(x\right)\left(x\right)=x^2=1.0\times {\mathrm{x}}^{\mathrm{-xiv}}$$

And then:

$$x=[{\text{H}}_3{\text{O}}^{\text{+}}]=[{\text{OH}}^{\text{−}}]=\sqrt{\mathrm{i.0}\times {10}^{\mathrm{-14}}}=1.0\times {10}^{\mathrm{-7}}M$$

The hydronium ion concentration and the hydroxide ion concentration are the same, one.0 $\times$ 10⁻⁷ Grand.

Check Your Learning

The ion product of water at 80 °C is two.four $\times$ x⁻¹³. What are the concentrations of hydronium and hydroxide ions in pure h2o at 80 °C?

Answer:

[H₃O⁺] = [OH⁻] = iv.9 $\times$ x⁻⁷ Thou

Example 14.2

The Inverse Relation betwixt [H₃O⁺] and [OH⁻]

A solution of an acid in water has a hydronium ion concentration of 2.0 $\times$ 10⁻⁶ M. What is the concentration of hydroxide ion at 25 °C?

Solution

Use the value of the ion-production abiding for water at 25 °C

$$\mathrm{two}{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{H}}_{\mathrm{three}}{\text{O}}^{\text{+}}(aq)+{\text{OH}}^{\text{−}}(aq){\mathrm{Chiliad}}_{\text{due west}}=[{\text{H}}_3{\text{O}}^{\text{+}}][{\text{OH}}^{\text{−}}]=1.0\times {10}^{\mathrm{-14}}$$

to calculate the missing equilibrium concentration.

Rearrangement of the K _westward expression shows that [OH⁻] is inversely proportional to [H_threeO⁺]:

$$[{\text{OH}}^{\text{−}}]=\frac{K_{\text{west}}}{[{\text{H}}_{\mathrm{iii}}{\text{O}}^{\text{+}}]}=\frac{\mathrm{i.0}\times {10}^{\mathrm{-14}}}{2.0\times {10}^{\mathrm{-six}}}=5.0\times {\mathrm{x}}^{\mathrm{-ix}}$$

Compared with pure h2o, a solution of acid exhibits a higher concentration of hydronium ions (due to ionization of the acid) and a proportionally lower concentration of hydroxide ions. This may exist explained via Le Châtelier's principle equally a left shift in the water autoionization equilibrium resulting from the stress of increased hydronium ion concentration.

Substituting the ion concentrations into the K _w expression confirms this calculation, resulting in the expected value:

$${\mathrm{One\; thousand}}_{\text{w}}=[{\text{H}}_3{\text{O}}^{\text{+}}][{\text{OH}}^{\text{−}}]=(2.0\times {\mathrm{ten}}^{\mathrm{-6}})(5.0\times {\mathrm{ten}}^{\mathrm{-9}})=1.0\times {10}^{\mathrm{-14}}$$

Check Your Learning

What is the hydronium ion concentration in an aqueous solution with a hydroxide ion concentration of 0.001 M at 25 °C?

Answer:

[H_threeO⁺] = 1 $\times$ ten^−xi Yard

Instance 14.3

Representing the Acrid-Base of operations Behavior of an Amphoteric Substance

Write carve up equations representing the reaction of ${\text{HSO}}_{\mathrm{iii}}{}^{\text{−}}$

(a) every bit an acid with OH⁻

(b) as a base with HI

Solution

(a) ${\text{HSO}}_3{}^{\text{−}}(aq)+{\text{OH}}^{\text{−}}(aq)\rightleftharpoons {\text{SO}}_3{}^{\text{2−}}(aq)+{\text{H}}_2\text{O}(l)$

(b) ${\text{HSO}}_3{}^{\text{−}}(aq)+\text{HI}(aq)\rightleftharpoons {\text{H}}_2{\text{SO}}_3(aq)+{\text{I}}^{\text{−}}(aq)$

Check Your Learning

Write dissever equations representing the reaction of ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}$

(a) as a base of operations with HBr

(b) as an acrid with OH⁻

Answer:

(a) ${\text{H}}_{\mathrm{two}}{\text{PO}}_{\mathrm{four}}{}^{\text{−}}(aq)+\text{HBr}(aq)\rightleftharpoons {\text{H}}_3{\text{PO}}_{\mathrm{iv}}(aq)+{\text{Br}}^{\text{−}}(aq);$ (b) ${\text{H}}_2{\text{PO}}_4{}^{\text{−}}(aq)+{\text{OH}}^{\text{−}}(aq)\rightleftharpoons {\text{HPO}}_4{}^{\text{two−}}(aq)+{\text{H}}_{\mathrm{two}}\text{O}(l)$

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Source: https://openstax.org/books/chemistry-2e/pages/14-1-bronsted-lowry-acids-and-bases

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