# 8.25: Brønsted-Lowry Acids and Bases: Autoionization of Water


Learning Objectives
• Write a Brønsted-Lowry acid/base equation that represents the autoionization of water.

Sections 8.17 through 8.24 explored the qualitative applications of Brønsted-Lowry acid/base reactivity.  As stated previously, strong acids and bases completely separate to form charged particles in solution.  In contrast, weak acids and bases undergo both "forward" dissociation and "reverse" recombination processes when dissolved in a solvent.  In order to quantitatively determine the extent to which these solutes ionize, chemists derived two equations that can be used to calculate the amounts of acid and base that are present in a solution.  These mathematical statements will be discussed and applied in the remaining sections of this chapter.  In order to establish a standard, or reference, value to which all subsequently-collected quantitative data could be compared, chemists initially studied a "solution" of "pure" liquid water, H2O.  Because no acidic or basic solutes were dissolved in this "solution," which was, therefore, classified as "neutral," scientists expected to detect only a single chemical, water, H2O.  Instead, these researchers discovered that very small amounts of acid and base were present in all of the water samples that were tested.

Recall that water, H2O, is amphoteric and, therefore, can act as both a Brønsted-Lowry acid, which is defined as a proton, H+1, donor in solution, and a Brønsted-Lowry base, which, by definition, accepts protons, H+1, that are created in solution.  A corresponding Brønsted-Lowry acid/base reaction, in which a proton, H+1, is transferred from one reactant water molecule to another, can be symbolically-represented in a Brønsted-Lowry acid/base equation, which will be described in the following paragraphs.

$$\ce{H_2O}$$ $$\left( l \right)$$$$\ce{H_2O}$$ $$\left( l \right)$$

Based on the information that is presented in Section 8.23, liquid water, H2O, is not classified as a strong acid and, therefore, is categorized as a weak acid, "by default."  Water, H2O, is also not one of the eight strong bases and, consequently, is classified as a weak base, "by default."  Therefore, because the acidic and basic reactants are both weak, an equilibrium arrow should be incorporated into the acid/base equation that is being developed, as shown below.

$$\ce{H_2O}$$ $$\left( l \right)$$$$\ce{H_2O}$$ $$\left( l \right)$$ $$\longrightleftharpoons$$

As explained in Sections 8.20, 8.21, and 8.22, the products that are generated during a Brønsted-Lowry acid/base reaction are the conjugates of the reactants that were initially present.  By definition, the chemical formulas of conjugate particles must differ by exactly and only one proton, H+1, and should otherwise be identical to one another.  Since a Brønsted-Lowry acid donates, or loses, a proton, H+1, during an acid/base reaction, the conjugate base that is formed as a result of this transfer contains one less proton than the acid from which it was generated.  Therefore, because water, H2O, contains two hydrogens, H, and one oxygen, O, and bears a net neutral charge, the loss of a proton, H+1, from this acid generates a particle that is comprised of one hydrogen, H, and one oxygen, O, and bears a net –1 charge.  Therefore, the conjugate base of water, H2O, is the hydroxide ion, OH–1.  Because all aqueous solutions that are prepared using basic solutes contain hydroxide ions, OH–1, the presence of this particle in a "solution" of "pure" water is chemically-significant.  Finally, because the "solution" that is being investigated was prepared using water, these hydroxide ions, OH–1, are produced in the aqueous state of matter, by definition.

$$\ce{H_2O}$$ $$\left( l \right)$$$$\ce{H_2O}$$ $$\left( l \right)$$ $$\longrightleftharpoons$$ $$\ce{OH^{–1}}$$ $$\left( aq \right)$$

Furthermore, because a Brønsted-Lowry base accepts, or gains, a proton, H+1, during an acid/base reaction, the conjugate acid that is formed as a result of this transfer contains one more proton than the base from which it was generated.  Therefore, because water, H2O, contains two hydrogens, H, and one oxygen, O, and bears a net neutral charge, the gain of a proton, H+1, by this base generates a particle that is comprised of three hydrogens, H, and one oxygen, O, and bears a net +1 charge.  Therefore, the conjugate acid of water, H2O, is the hydronium ionH3O+1.  As stated previously, the hydronium ionH3O+1, is always generated in aqueous solutions that contain acidic solutes, and, therefore, the presence of this particle in a "solution" of "pure" water is chemically-significant.  Finally, because the "solution" that is being investigated was prepared using water, these hydronium ionsH3O+1, are generated in the aqueous state of matter, by definition.

$$\ce{H_2O}$$ $$\left( l \right)$$$$\ce{H_2O}$$ $$\left( l \right)$$ $$\longrightleftharpoons$$ $$\ce{OH^{–1}}$$ $$\left( aq \right)$$ + $$\ce{H_3O^{+1}}$$ $$\left( aq \right)$$

Because all of the components in the final equation that is shown above are balanced, this equation is the chemically-correct representation of the Brønsted-Lowry acid/base reaction that occurs between two watemolecules.  Since this process spontaneously occurs in all water samples, this transformation is known as the autoionization of water.

8.25: Brønsted-Lowry Acids and Bases: Autoionization of Water is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.