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15.1: Bronsted-Lowry Acids and Bases

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Bronsted-Lowry Acids and Bases

15.1: Bronsted-Lowry Acids and Bases

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 CO2), 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.

Brønsted-Lowry Acids and Bases

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, 1H.) 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 H+, 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 H+, it is converted to its conjugate acid. The reaction between water and ammonia illustrates this idea as shown below.


In the forward direction, water acts as an acid by donating a proton to ammonia and subsequently becoming a hydroxide ion, OH, the conjugate base of water. The ammonia acts as a base in accepting this proton, becoming an ammonium ion, NH4+, 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.

Strong acids and bases dissociate completely in a solution. Their conjugate acids and bases are extremely weak and can not donate or accept the protons, respectively, to carry out the reverse reaction; therefore, reactions involving strong acids and bases essentially go to completion when in an aqueous solution. On the other hand, weak acids and bases partially dissociate in solutions and produce weak conjugate bases and acids, respectively. These weak conjugate acids or bases can carry out the reverse reaction, and thus reactions of weak acid and base reach an equilibrium depending upon the relative strengths of the weak acids and bases. To summarize, a stronger acid will produce the equally weaker conjugate base whereas a stronger base will produce the equally weaker conjugate acid and vice versa. Table 1 depicts the relation between different conjugate acid-base pairs.

Strong Acid Very Weak Conjugate Base
HCl Cl
Weak Acid Weak Conjugate Base
NH4+ NH3
Very Weak Acid Strong Conjugate Base
OH O2−

Table 1: Relative strength of a few conjugate acid-base pairs.

This text is adapted from Openstax, Chemistry 2e, Section 14.4 Brønsted-Lowry Acid and Bases.

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