Brønsted–Lowry Acids and Bases

Introduction: why do acids and bases matter? 🌟

students, imagine squeezing lemon juice on food, adding antacid tablets to calm stomach acid, or using fertilizers that change soil pH. All of these connect to acid-base chemistry, one of the most important parts of reactivity in chemistry. In IB Chemistry SL, the Brønsted–Lowry model helps explain how acids and bases react by focusing on proton transfer.

By the end of this lesson, you should be able to:

explain the Brønsted–Lowry definition of acids and bases,
identify acids, bases, conjugate acid-base pairs, and amphiprotic species,
write and interpret proton-transfer equations,
connect these ideas to reaction mechanisms in chemical change,
use examples and evidence to explain acid-base behavior in real situations.

This topic is important because it helps you see chemical change as a transfer process, not just a list of substances. In many reactions, what really moves is a proton, $\mathrm{H^+}$, and that movement changes the identity and properties of the molecules involved.

Brønsted–Lowry definition: proton transfer in action

The Brønsted–Lowry theory defines an acid as a proton donor and a base as a proton acceptor. A proton is the nucleus of a hydrogen atom, so in chemistry we often describe acid-base reactions as the transfer of $\mathrm{H^+}$ from one species to another.

A general reaction can be shown as:

$$\mathrm{HA + B \rightleftharpoons A^- + BH^+}$$

In this equation:

$\mathrm{HA}$ is the acid because it donates $\mathrm{H^+}$,
$\mathrm{B}$ is the base because it accepts $\mathrm{H^+}$,
$\mathrm{A^-}$ is the conjugate base,
$\mathrm{BH^+}$ is the conjugate acid.

This is a mechanism of chemical change because the key event is the movement of a proton from one particle to another. The molecules are not simply “mixing”; their identities change through electron pair interactions that make proton transfer possible.

Example: hydrochloric acid in water.

$$\mathrm{HCl + H_2O \rightarrow H_3O^+ + Cl^-}$$

Here, $\mathrm{HCl}$ is the acid because it donates $\mathrm{H^+}$, and $\mathrm{H_2O}$ is the base because it accepts $\mathrm{H^+}$. The products are the conjugate base $\mathrm{Cl^-}$ and the conjugate acid $\mathrm{H_3O^+}$.

Conjugate acid-base pairs: linked by one proton 🔗

Every Brønsted–Lowry acid-base reaction creates two conjugate pairs. A conjugate acid-base pair differs by exactly one proton, $\mathrm{H^+}$.

For the reaction

$$\mathrm{NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-}$$

there are two conjugate pairs:

$\mathrm{NH_3}$ and $\mathrm{NH_4^+}$,
$\mathrm{H_2O}$ and $\mathrm{OH^-}$.

Here, $\mathrm{NH_3}$ acts as the base because it accepts $\mathrm{H^+}$ from water. Water acts as the acid because it donates $\mathrm{H^+}$. The reaction is reversible, so both forward and reverse proton transfers can occur.

A useful rule is this: the stronger the acid, the weaker its conjugate base. For example, $\mathrm{HCl}$ is a strong acid, so its conjugate base $\mathrm{Cl^-}$ is very weak and does not easily accept a proton. In contrast, a weak acid like ethanoic acid has a conjugate base that can still accept a proton to some extent.

This relationship helps explain why some acid-base reactions go nearly to completion while others establish an equilibrium. In chemistry, the position of equilibrium reflects the relative strengths of the acids and bases involved.

Amphiprotic species: can act as either acid or base 🔄

Some substances can behave as either a Brønsted–Lowry acid or a Brønsted–Lowry base depending on what they react with. These are called amphiprotic species.

Water is the classic example. It can donate a proton:

$$\mathrm{H_2O + NH_3 \rightleftharpoons OH^- + NH_4^+}$$

or accept a proton:

$$\mathrm{H_2O + HCl \rightarrow H_3O^+ + Cl^-}$$

Other amphiprotic species include $\mathrm{HCO_3^-}$ and $\mathrm{HSO_4^-}$.

Take $\mathrm{HCO_3^-}$ as an example. It can donate a proton to form $\mathrm{CO_3^{2-}}$:

$$\mathrm{HCO_3^- \rightarrow CO_3^{2-} + H^+}$$

or accept a proton to form $\mathrm{H_2CO_3}$:

$$\mathrm{HCO_3^- + H^+ \rightarrow H_2CO_3}$$

In exam questions, the key is to look at the reaction partner. If a species gives away $\mathrm{H^+}$, it is acting as an acid. If it takes $\mathrm{H^+}$, it is acting as a base.

Strong and weak acids and bases: not all reactions behave the same ⚖️

In Brønsted–Lowry chemistry, strength refers to the extent of proton donation or acceptance in water, not the concentration of the solution. A strong acid dissociates almost completely, while a weak acid only partially dissociates.

Examples of strong acids commonly studied include $\mathrm{HCl}$, $\mathrm{HNO_3}$, and $\mathrm{H_2SO_4}$ for its first dissociation. A weak acid example is ethanoic acid, $\mathrm{CH_3COOH}$.

For ethanoic acid in water:

$$\mathrm{CH_3COOH + H_2O \rightleftharpoons CH_3COO^- + H_3O^+}$$

Because this reaction is reversible, an equilibrium is formed. That means at any moment, some molecules are donating protons while others are accepting them back.

Strong and weak bases work similarly. Sodium hydroxide is a strong base because it dissociates to release $\mathrm{OH^-}$ in water, and ammonia is a weak base because it accepts protons only partially:

$$\mathrm{NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-}$$

A common IB idea is that a strong acid has a weak conjugate base, and a strong base has a weak conjugate acid. This helps you predict direction and extent of reaction.

How to identify acid-base reactions in IB questions 🧠

To apply Brønsted–Lowry reasoning, students, follow these steps:

Find the species with a hydrogen atom that can be transferred.
Check which species loses $\mathrm{H^+}$ and which gains $\mathrm{H^+}$.
Label the acid, base, conjugate acid, and conjugate base.
Compare reactants and products to identify conjugate pairs.

Example:

$$\mathrm{HCO_3^- + H_2O \rightleftharpoons H_2CO_3 + OH^-}$$

$\mathrm{HCO_3^-}$ is the base because it accepts $\mathrm{H^+}$,
$\mathrm{H_2O}$ is the acid because it donates $\mathrm{H^+}$,
$\mathrm{H_2CO_3}$ is the conjugate acid,
$\mathrm{OH^-}$ is the conjugate base.

In another reaction:

$$\mathrm{HCO_3^- + H_2O \rightleftharpoons CO_3^{2-} + H_3O^+}$$

$\mathrm{HCO_3^-}$ now acts as an acid, showing that amphiprotic species must be identified from the actual reaction, not from a fixed label.

A good exam habit is to write proton transfer explicitly when possible. That makes the mechanism clear and reduces errors in naming species.

Connection to broader Reactivity 3: mechanisms of chemical change ⚗️

This lesson fits into Reactivity 3 because it explains a mechanism of change at the particle level. In acid-base reactions, the main change is proton transfer. This is similar in spirit to other mechanistic chemistry topics where you explain how substances react, not only what the products are.

Brønsted–Lowry ideas also connect to:

redox and electrochemistry, because some reactions involve both proton transfer and electron transfer in different contexts,
organic reaction pathways, where acids and bases often act as catalysts or reactants,
equilibrium, because many acid-base reactions are reversible and established as dynamic equilibria.

For example, in biological systems, amino acids can gain or lose protons depending on pH. In the stomach, acid helps digestion by changing the structure of proteins. In soil, acidity affects nutrient availability for plants. These are all real-world examples of proton-transfer chemistry shaping reactivity and behavior.

Conclusion: the big idea to remember ✅

Brønsted–Lowry acid-base chemistry is about proton transfer. An acid donates $\mathrm{H^+}$, a base accepts $\mathrm{H^+}$, and every reaction forms conjugate pairs that differ by one proton. Some species, like water and bicarbonate, are amphiprotic and can play either role depending on the reaction.

students, if you can identify who gives the proton, who receives it, and what the conjugate pairs are, you can analyze most IB SL acid-base questions confidently. This model is a powerful tool because it turns a familiar topic into a clear explanation of chemical mechanism.

Study Notes

A Brønsted–Lowry acid is a proton donor, $\mathrm{H^+}$ donor.
A Brønsted–Lowry base is a proton acceptor, $\mathrm{H^+}$ acceptor.
Conjugate acid-base pairs differ by exactly one proton.
Strong acids have weak conjugate bases; strong bases have weak conjugate acids.
Water is amphiprotic, so it can act as either an acid or a base.
In equations, identify which species loses $\mathrm{H^+}$ and which gains $\mathrm{H^+}$.
Many acid-base reactions are reversible and establish equilibrium.
This topic is part of Reactivity 3 because it explains a mechanism of chemical change through proton transfer.
Real-world examples include stomach acid, antacids, blood buffering, and soil chemistry.