2.7 Acids and Bases: The Brønsted–Lowry Definition

Brønsted–Lowry acid–base chemistry centers on one simple event: a proton ($H^+$) moves from one molecule to another. Every reaction you'll see in this framework is just a proton transfer, and learning to track where that proton goes is essential for predicting reactivity and understanding equilibrium throughout organic chemistry.

Brønsted–Lowry Acid–Base Reactions

Components of Brønsted–Lowry Reactions

A Brønsted–Lowry acid is a proton donor; a Brønsted–Lowry base is a proton acceptor. When the proton transfers, two new species form:

• Conjugate base: what the acid becomes after it loses $H^+$. It has one fewer proton and is one unit more negative than the original acid. For example, $HCl$ donates its proton and becomes $Cl^-$.
• Conjugate acid: what the base becomes after it gains $H^+$. It has one more proton and is one unit more positive than the original base. For example, $NH_3$ accepts a proton and becomes $NH_4^+$.

Every Brønsted–Lowry reaction therefore contains two conjugate acid–base pairs. Being able to identify both pairs in a reaction is a skill you'll use constantly.

Amphoteric Nature of Water

Water is amphoteric, meaning it can act as either an acid or a base depending on what it reacts with.

• Water as a base (reacting with a stronger acid): $HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-$. Here $H_2O$ accepts a proton from $HCl$, forming the conjugate acid $H_3O^+$ (hydronium) and the conjugate base $Cl^-$.
• Water as an acid (reacting with a stronger base): $H_2O + NH_3 \rightleftharpoons OH^- + NH_4^+$. Here $H_2O$ donates a proton to $NH_3$, forming the conjugate base $OH^-$ and the conjugate acid $NH_4^+$.

The key idea: water adjusts its role based on the relative strength of the other species in the reaction. A stronger acid than water will force water to act as a base, and a stronger base than water will force water to act as an acid.

Proton Transfer in Acid–Base Reactions

Tracking proton transfer follows a predictable pattern:

1. Identify the acid (the species with the proton to donate) and the base (the species that will accept it).
2. Move $H^+$ from the acid to the base.
3. Write the products: the acid minus $H^+$ is the conjugate base; the base plus $H^+$ is the conjugate acid.

For a generic reaction: $HA + B \rightleftharpoons A^- + HB^+$

The double arrow ($\rightleftharpoons$) tells you the reaction reaches equilibrium rather than going to completion. Which side is favored? That depends on relative acid and base strengths. The proton will preferentially transfer from the stronger acid to the stronger base, and equilibrium will favor the side with the weaker acid and weaker base. This is a pattern worth memorizing: equilibrium favors formation of the weaker acid and weaker base.

Extended Acid–Base Concepts

The Brønsted–Lowry definition focuses entirely on proton transfer. You should know that the Lewis acid–base theory broadens the picture by defining acids as electron-pair acceptors and bases as electron-pair donors. Lewis theory covers reactions where no proton transfer occurs at all (such as $BF_3$ accepting a lone pair from $NH_3$). You'll encounter Lewis acids and bases frequently later in the course, especially in reaction mechanisms, but for now the Brønsted–Lowry framework is the foundation to master.