The study of acids and bases is fundamental to chemistry, describing how substances behave when mixed. Early definitions were limited, often requiring the presence of water for classification. In 1923, Danish chemist Johannes Nicolaus Brønsted and English chemist Thomas Martin Lowry independently proposed a new, more expansive theory. The Brønsted-Lowry theory shifted the focus from resulting ions to the actual chemical action: the transfer of a subatomic particle. This definition provided a comprehensive framework applicable across various chemical environments, including those that do not involve water.
Defining Proton Acceptors and Donors
The Brønsted-Lowry model defines acids and bases solely by their actions during a reaction involving proton transfer. A Brønsted-Lowry acid functions as a proton donor, giving away a hydrogen ion (\(H^+\)). Conversely, a Brønsted-Lowry base acts as a proton acceptor, taking in the hydrogen ion donated by the acid. A proton is a hydrogen atom that has lost its single electron, represented as \(H^+\).
The \(H^+\) ion is the mechanism of the acid-base reaction. To illustrate, consider the reaction between hydrochloric acid (\(HCl\)) and ammonia (\(NH_3\)). \(HCl\) acts as the acid, donating \(H^+\). Ammonia acts as the base, accepting the \(H^+\) to form the ammonium ion (\(NH_4^+\)).
A base’s ability to accept a proton is often due to a lone pair of electrons, which form a new bond with the incoming hydrogen ion. This definition is flexible because it only requires proton movement between the two reactant species, explaining acid-base behavior even in non-aqueous environments.
Understanding Conjugate Acid-Base Pairs
The exchange of a proton between a Brønsted-Lowry acid and a base creates a new acid and a new base, known as a conjugate pair. The original acid, after donating its proton, becomes the conjugate base. The original base, after accepting the proton, becomes the conjugate acid. This relationship is always reversible and represented by a dynamic equilibrium.
The general reaction is written as \(HA + B \rightleftharpoons A^- + HB^+\). Here, \(HA\) and \(A^-\) form one conjugate pair, and \(B\) and \(HB^+\) form the other. \(A^-\) is the conjugate base (the original acid minus a proton), and \(HB^+\) is the conjugate acid (the original base plus a proton).
A primary concept is the inverse relationship between the strength of an acid or base and its conjugate. A strong acid, such as \(HCl\), readily gives up its proton. This means its resulting conjugate base (e.g., the chloride ion, \(Cl^-\)) has a very low tendency to re-accept a proton. Therefore, a strong acid always yields a very weak conjugate base, and a strong base yields a very weak conjugate acid. This inverse proportionality is quantitatively linked: the product of the acid dissociation constant (\(K_a\)) and the base dissociation constant (\(K_b\)) for a given conjugate pair equals the ion product of water (\(K_w\)).
Why the Bronsted-Lowry Model is Necessary
The Brønsted-Lowry model became necessary because the previous standard, the Arrhenius definition, was too narrow. The Arrhenius model defined acids as substances producing hydrogen ions (\(H^+\)) and bases as substances producing hydroxide ions (\(OH^-\)) only when dissolved in water. This framework meant that reactions occurring outside of an aqueous solution could not be classified.
A major limitation of the Arrhenius theory was its inability to classify common substances like ammonia (\(NH_3\)) as a base. Since ammonia does not contain \(OH^-\) in its formula, it did not fit the Arrhenius definition, despite exhibiting basic properties. The Brønsted-Lowry definition solves this by recognizing that ammonia acts as a proton acceptor in any environment, classifying it as a base without requiring water or hydroxide ions.
The Brønsted-Lowry framework also introduced the concept of amphiprotic substances, highlighting molecular versatility. An amphiprotic substance can act as either a proton donor (acid) or a proton acceptor (base), depending on the chemical it reacts with. Water (\(H_2O\)) is the most frequently cited example, acting as a base when reacting with \(HCl\) and acting as an acid when reacting with \(NH_3\).