DAT Organic Chemistry: Bronsted Acids, Bases, and Conjugates

DAT High-Yield Concept

Brønsted-Lowry Definition: A Brønsted acid is a proton ( $\ce{H^+}$ ) donor, while a Brønsted base is a proton ( $\ce{H^+}$ ) acceptor.
Conjugate Pairs: Every acid-base reaction forms two conjugate pairs that differ by exactly one proton ( $\ce{H^+}$ ). Removing a proton yields the conjugate base; adding a proton yields the conjugate acid.
Inverse Strength Relationship: The stronger an acid, the weaker (more stable and less reactive) its conjugate base.
Amphiprotic Species: Molecules or ions that can either donate or accept a proton depending on their chemical environment (e.g., $\ce{H2O}$ , $\ce{HCO3^-}$ ).

1. Brønsted-Lowry Acids and Bases

In organic chemistry, acid-base reactions describe proton transfer dynamics. Under the Brønsted-Lowry model, a Brønsted acid donates a proton ( $\ce{H^+}$ ) to a receiver, whereas a Brønsted base accepts a proton using a lone pair of electrons or a $\pi$ bond. Unlike Arrhenius theory—which requires an aqueous environment—Brønsted-Lowry theory applies to organic solvents and gas-phase transformations.

graphic showing the relationship between three acid-base theories (Source: Tem5psu · CC BY-SA 3.0 · Wikimedia Commons)

During a Brønsted proton transfer, curved arrows originate at the electron source (a lone pair on the base) and attack the acidic hydrogen. Simultaneously, the electron pair from the $\ce{H-A}$ bond shifts onto atom $A$ , forming the conjugate base.

Key Distinction: Brønsted acidity specifically involves proton (

\ce{H^+}

) transfer. Lewis acidity involves electron-pair acceptance and encompasses broader electrophile-nucleophile additions (covered in depth in its own lesson).

2. Drawing Conjugate Acids and Conjugate Bases

A conjugate acid-base pair consists of two chemical species that differ by precisely one proton ( $\ce{H^+}$ ).

To draw the conjugate base of a given acid: remove one $\ce{H^+}$ and decrease the net formal charge by 1.
To draw the conjugate acid of a given base: add one $\ce{H^+}$ and increase the net formal charge by 1.

A conjugate pair differs by exactly one proton — Conjugate is a relationship, not a property. Acetic acid gives up one proton to become acetate, so those two are a conjugate acid-base pair; ammonia takes one proton on to become ammonium, so those two are also a pair. In both cases the partners differ by exactly one H+ and by exactly one unit of charge, and by nothing else. If two species differ by two protons, or by a proton plus a rearranged skeleton, they are not conjugates.

When acetic acid dissociates in water, acetic acid donates a proton to water, producing acetate and hydronium:

$\ce{CH3COOH + H2O <=> CH3COO^- + H3O^+}$

Ball-and-stick model of the dissociation of acetic acid to acetate. A water molecule is protonated to form a hydronium ion in the process. The acidic proton that is transferred from acetic acid to water is labelled in green. Image generated in Accelrys DS Visualizer . (Source: Ben Mills · Public domain · Wikimedia Commons)

Mnemonic: AD-BA — Acids Donate, Bases Accept. Adding a proton increases charge by +1; removing a proton decreases charge by -1!

Worked Example 1: Drawing Conjugates

Problem: Identify (a) the conjugate base of phenol ( $\ce{C6H5OH}$ ) and (b) the conjugate acid of methylamine ( $\ce{CH3NH2}$ ).

Reasoning:

For phenol ( $\ce{C6H5OH}$ ), remove the acidic O-H proton ( $\ce{H^+}$ ). Neutral charge $0 - 1 = -1$ . The conjugate base is phenoxide, $\ce{C6H5O^-}$ .
For methylamine ( $\ce{CH3NH2}$ ), add $\ce{H^+}$ to the basic nitrogen lone pair. Neutral charge $0 + 1 = +1$ . The conjugate acid is methylammonium, $\ce{CH3NH3^+}$ .

DAT Exam Trap: Do not remove non-acidic carbon-bound protons (e.g., removing a proton from the methyl group of acetic acid to yield $\ce{^-CH2COOH}$ ). Always target the most acidic heteroatom-bound hydrogen!

3. Amphiprotic Species

An amphiprotic species is a molecule or ion that can act as either a Brønsted acid (proton donor) or a Brønsted base (proton acceptor). To be amphiprotic, a structure must possess at least one transferable hydrogen atom and at least one lone pair of electrons.

Common Amphiprotic Species: Water (

\ce{H2O}

), bicarbonate (

\ce{HCO3^-}

), hydrogen sulfate (

\ce{HSO4^-}

), and amino acid zwitterions.

Worked Example 2: Dual Reactivity of Bicarbonate

Problem: Write balanced Brønsted-Lowry equations demonstrating the amphiprotic nature of the bicarbonate ion ( $\ce{HCO3^-}$ ) in aqueous solution.

Reasoning:

As a Base: In the presence of hydronium, $\ce{HCO3^-}$ accepts a proton to yield carbonic acid: $\ce{HCO3^- + H3O^+ <=> H2CO3 + H2O}$
As an Acid: In the presence of hydroxide, $\ce{HCO3^-}$ donates a proton to yield carbonate: $\ce{HCO3^- + OH^- <=> CO3^{2-} + H2O}$

DAT Exam Trap: All amphiprotic species are amphoteric (can react as acid or base), but not all amphoteric species are amphiprotic! For example, metal oxides like $\ce{Al2O3}$ are amphoteric via Lewis pathways, but contain no protons to transfer.

4. Inverse Strength Relationship of Conjugate Pairs

A foundational principle of organic reactivity is the inverse strength relationship: the stronger an acid, the weaker (more stable) its conjugate base. Conversely, a weak acid yields a strong, reactive conjugate base.

Structural factors that stabilize a conjugate base—such as high electronegativity, larger atomic radius, resonance delocalization, inductive electron-withdrawal, and high s-character hybridization—reduce its basicity. Consequently, the parent acid becomes significantly more acidic (quantitative treatment via $K_a$ and $\mathrm{p}K_a$ is covered in depth in its own lesson).

Worked Example 3: Predicting Relative Base Strength

Problem: Compare the relative basicity of ethoxide ( $\ce{CH3CH2O^-}$ ) and acetate ( $\ce{CH3COO^-}$ ).

Reasoning:

Identify parent conjugate acids: ethanol ( $\ce{CH3CH2OH}$ ) and acetic acid ( $\ce{CH3COOH}$ ).
Evaluate conjugate base stability: In acetate ( $\ce{CH3COO^-}$ ), negative charge is delocalized over two oxygen atoms via resonance. In ethoxide ( $\ce{CH3CH2O^-}$ ), negative charge is localized on a single oxygen atom.
Apply inverse strength: Because acetate is a far more stable (weaker) conjugate base, acetic acid is a stronger acid than ethanol ( $\ce{CH3COOH}$ > $\ce{CH3CH2OH}$ ). Therefore, ethoxide is the stronger base ( $\ce{CH3CH2O^-}$ > $\ce{CH3COO^-}$ ).

DAT Exam Trap: Do not mistake conjugate base stability for base strength. Increased conjugate base stability means lower basicity and higher parent acidity!

Summary & Comparison Table

Theory / Concept	Acid Definition	Base Definition	Key Characteristic
Arrhenius	Produces $\ce{H^+}$ ( $\ce{H3O^+}$ ) in water	Produces $\ce{OH^-}$ in water	Restricted exclusively to aqueous solutions.
Brønsted-Lowry	Proton ( $\ce{H^+}$ ) donor	Proton ( $\ce{H^+}$ ) acceptor	Requires proton transfer; forms conjugate pairs.
Amphiprotic Species	Donates $\ce{H^+}$ to strong bases	Accepts $\ce{H^+}$ from strong acids	Must possess both a transferable $\ce{H}$ and lone pairs.
Conjugate Inverse Strength	Stronger acid $\rightarrow$ weaker conjugate base	Stronger base $\rightarrow$ weaker conjugate acid	Conjugate stability dictates relative acidity/basicity.

Critical Thinking Questions

The dihydrogen phosphate ion ( $\ce{H2PO4^-}$ ) is amphiprotic. Write balanced chemical equations for its reaction with (a) $\ce{H3O^+}$ and (b) $\ce{OH^-}$ .
Explain why the conjugate base of trifluoroacetic acid ( $\ce{CF3COOH}$ ) is much weaker than the conjugate base of acetic acid ( $\ce{CH3COOH}$ ). Which parent acid is stronger?
Given the series $\ce{NH3}$ , $\ce{H2O}$ , and $\ce{HF}$ , arrange their conjugate bases in order of increasing basicity and justify using electronegativity trends.

References

Sources returned by Google Search grounding during lesson verification.

2.6: Acids and Bases - The Brønsted-Lowry Definition - Chemistry LibreTexts — chem.libretexts.org
Acid-Base Organic Chemistry for the DAT - Shemmassian Academic Consulting — shemmassianconsulting.com

Bronsted Acids, Bases, and Conjugates

Summary

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