15.1:
Bronsted-Lowry Acids and Bases
An Arrhenius acid is a substance that produces a hydrogen ion when dissolved in water, and an Arrhenius base is a substance that produces an OH−, or hydroxide ion.
Hydrochloric acid is an Arrhenius acid as it dissociates into a hydrogen ion and a chloride ion when dissolved in water. Sodium hydroxide is an Arrhenius base as it dissociates into a sodium ion and a hydroxide ion when dissolved in water.
However, this definition cannot be used to describe acids and bases that are not in an aqueous solution or bases that do not contain hydroxide ions.
A broader definition by Brønsted and Lowry defines an acid as a hydrogen ion or proton donor, whereas a base is a proton acceptor.
When hydrochloric acid is dissolved in water, it acts as an acid by donating a proton to water, producing a hydronium ion and a chloride ion. When ammonia is dissolved in water, it acts as a base and accepts a proton from water, producing an ammonium ion and a hydroxide ion.
A Brønsted-Lowry acid will always react with a Brønsted-Lowry base and vice versa. When an acid, for example, acetic acid, donates its proton, water acts as a base and accepts the proton. Acetic acid is converted into a conjugate base, acetate, and water is converted into a conjugate acid, a hydronium ion.
Acids and bases that differ from each other due to the transfer of a proton are called conjugate acid-base pairs. In the reverse reaction, the conjugate acid, hydronium, acts as a proton donor and the conjugate base, acetate, will accept a proton.
The strength of an acid is determined by its ability to donate a proton whereas the strength of a base is determined by its ability to accept a proton. A stronger acid is more likely to donate a proton than a weaker acid. Likewise, a stronger base is more likely to accept a proton than a weaker base.
The strength of an acid and its conjugate base are inversely related. A strong acid dissociates completely in solution and the resulting conjugate base is too weak to accept a proton. The same applies in the case of a strong base and its conjugate acid.
On the other hand, a weak acid dissociates partially in solution. The conjugate base of a weak acid is also relatively weak; therefore, a mixture of the undissociated weak acid and its weak conjugate base will be present in equilibrium. The same phenomenon occurs in the case of a weak base and its weak conjugate acid.
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− |
| HNO3 | NO3− |
| Weak Acid | Weak Conjugate Base |
| HF | F− |
| NH4+ | NH3 |
| Very Weak Acid | Strong Conjugate Base |
| OH− | O2− |
| CH4 | CH3− |
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.
Suggested Reading
- Story, David A. "Bench-to-bedside review: A brief history of clinical acid–base." Critical Care 8, no. 4 (2004): 253.