Brønsted-Lowry Acids and Bases: Concepts, Classification, and Examples

Brønsted-Lowry Acids and Bases

Introduction to Brønsted-Lowry Theory

The Brønsted-Lowry theory expands the definition of acids and bases beyond the Arrhenius concept by focusing on proton (H+) transfer. This theory is fundamental to understanding acid-base reactions in aqueous and non-aqueous solutions.

• Brønsted-Lowry Acid: A species that donates a proton (H+) to another substance.

• Brønsted-Lowry Base: A species that accepts a proton (H+) using its lone pair of electrons.

Key Point: The hydrogen ion (H+) is essentially a proton, as a hydrogen atom consists of one proton and one electron. When the electron is lost, only the proton remains.

Visualizing Proton Transfer

• In a typical Brønsted-Lowry acid-base reaction, the acid donates a proton to the base, which accepts it via a lone pair of electrons.

• Example (generalized):

$\mathrm{HA} + \mathrm{B} \rightarrow \mathrm{A}^- + \mathrm{HB}^+$

• Here, HA is the acid (proton donor), and B is the base (proton acceptor).

Classification of Acids and Bases

Acids and bases can be classified according to different theories. The Brønsted-Lowry definition is broader than the Arrhenius definition but narrower than the Lewis definition.

Type	Acid Definition	Base Definition
Arrhenius	Produces H+ in water	Produces OH- in water
Brønsted-Lowry	Proton (H+) donor	Proton (H+) acceptor
Lewis	Electron pair acceptor	Electron pair donor

Examples of Brønsted-Lowry Acid-Base Reactions

• Example 1:

$\mathrm{HCl}_{(aq)} + \mathrm{OH}^-_{(aq)} \rightarrow \mathrm{Cl}^-_{(aq)} + \mathrm{H_2O}_{(l)}$

• HCl is the acid (donates H+), OH- is the base (accepts H+).

• Example 2:

$\mathrm{HCl}_{(aq)} + \mathrm{NH_3}_{(aq)} \rightarrow \mathrm{NH_4Cl}_{(aq)}$

• HCl is the acid, NH3 is the base.

Classifying Reactants in Equilibrium Equations

• In acid-base equilibrium reactions, identify which species donates and which accepts a proton.

Equation	Acid	Base
$\mathrm{HF} + \mathrm{SO_3^{2-}} \rightleftharpoons \mathrm{F^-} + \mathrm{HSO_3^-}$	HF	SO32-
$\mathrm{CO_3^{2-}} + \mathrm{CH_3COOH} \rightleftharpoons \mathrm{CH_3COO^-} + \mathrm{HCO_3^-}$	CH3COOH	CO32-
$\mathrm{H_3PO_4} + \mathrm{OCl^-} \rightleftharpoons \mathrm{H_2PO_4^-} + \mathrm{HOCl}$	H3PO4	OCl-
$\mathrm{HCO_3^-} + \mathrm{HSO_4^-} \rightleftharpoons \mathrm{SO_4^{2-}} + \mathrm{H_2CO_3}$	HSO4-	HCO3-

Monoprotic vs. Polyprotic Acids and Bases

Monoprotic Acids

Monoprotic acids are acids that can donate only one proton (H+) per molecule in an acid-base reaction.

• Examples: HCl, HNO3, HF

General dissociation equations:

• $\mathrm{HCl}_{(aq)} \rightarrow \mathrm{H^+} + \mathrm{Cl^-}$

• $\mathrm{HNO_3}_{(aq)} \rightarrow \mathrm{H^+} + \mathrm{NO_3^-}$

• $\mathrm{HF}_{(aq)} \rightarrow \mathrm{H^+} + \mathrm{F^-}$

Polyprotic Acids

Polyprotic acids can donate more than one proton per molecule. They dissociate in steps, each releasing one proton at a time.

• Diprotic acids: Can donate two protons (e.g., H2SO4, H2CO3).

• Triprotic acids: Can donate three protons (e.g., H3PO4).

Examples of stepwise dissociation:

• Sulfuric acid (H2SO4):

• $\mathrm{H_2SO_4} \rightarrow \mathrm{H^+} + \mathrm{HSO_4^-}$

• $\mathrm{HSO_4^-} \rightarrow \mathrm{H^+} + \mathrm{SO_4^{2-}}$

• Phosphoric acid (H3PO4):

• $\mathrm{H_3PO_4} \rightarrow \mathrm{H^+} + \mathrm{H_2PO_4^-}$

• $\mathrm{H_2PO_4^-} \rightarrow \mathrm{H^+} + \mathrm{HPO_4^{2-}}$

• $\mathrm{HPO_4^{2-}} \rightarrow \mathrm{H^+} + \mathrm{PO_4^{3-}}$

Organic Acids: Monoprotic and Polyprotic

Organic acids, such as carboxylic acids, are typically monoprotic if they contain a single –COOH group. Polyprotic organic acids contain more than one –COOH group.

Acid	Formula	Type	Dissociation Equation
Methanoic acid	HCOOH	Monoprotic	$\mathrm{HCOOH} \rightarrow \mathrm{H^+} + \mathrm{HCOO^-}$
Ethanoic acid	CH3COOH	Monoprotic	$\mathrm{CH_3COOH} \rightarrow \mathrm{H^+} + \mathrm{CH_3COO^-}$
Oxalic acid	HOOCCOOH	Diprotic	$\mathrm{HOOCCOOH} \rightarrow 2\mathrm{H^+} + \mathrm{OOCCOO^{2-}}$
Citric acid	C6H8O7	Triprotic	$\mathrm{C_6H_8O_7} \rightarrow 3\mathrm{H^+} + \mathrm{C_6H_5O_7^{3-}}$

Identifying Acids, Bases, and Alcohols

• Acids: Typically contain hydrogen that can be released as H+ (e.g., –COOH group in carboxylic acids).

• Alcohols: Contain the –OH group bonded to carbon, but do not release H+ as readily as acids.

• Rule: If a compound contains a –COOH group, it is likely a carboxylic acid. If it contains an –OH group attached to a saturated carbon, it is likely an alcohol.

Summary

• Brønsted-Lowry acids are proton donors; bases are proton acceptors.

• Monoprotic acids donate one proton per molecule; polyprotic acids donate more than one.

• Organic acids with one –COOH group are monoprotic; those with multiple –COOH groups are polyprotic.

Example Question: Which of the following can act as a Brønsted-Lowry base? (Answer: The species that can accept a proton, such as HCO3- in NaHCO3.)