What is the difference between an acid and a base?

An acid is a proton donor, while a base is a proton acceptor. This is the Brønsted-Lowry definition used to explain proton transfer reactions in IB Chemistry. When an acid and a base react, the acid transfers an H⁺ ion to the base, producing a new conjugate acid-base pair.

This distinction is more accurate than describing acids as sour and bases as slippery. It explains why substances such as ammonia can be bases even though they contain no hydroxide ions, and why water can behave as either an acid or a base. This article focuses on that central difference, with the conjugate pairs, pH relationships, examples, and exam reasoning needed to apply it correctly.

Acid vs base chemistry at a glance

The current IB Chemistry course places acids and bases within Reactivity 3.1: Proton transfer reactions. The official IB Chemistry subject brief identifies proton transfer as part of the course's reactivity framework.

The definitions apply to chemical species, not only to neutral compounds. Molecules such as HCl and NH₃, cations such as NH₄⁺, and anions such as HCO₃⁻ can all participate in acid-base reactions.

What makes a substance an acid?

According to the Brønsted-Lowry theory, an acid is a species capable of donating a proton, H⁺, to another species. This agrees with the IUPAC definition of a Brønsted acid.

Consider hydrogen chloride reacting with water:

HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)

HCl donates H⁺, so it acts as the acid. Water accepts the proton and therefore acts as the base. Free protons do not exist independently to any significant extent in water, so H⁺ is transferred to H₂O to form H₃O⁺, the hydronium ion.

IB equations and calculations often use H⁺(aq) as a convenient shorthand for hydrated hydrogen ions. You should nevertheless understand that H₃O⁺ represents the proton associated with water more explicitly.

Everyday acid examples

Common acidic substances include:

• Vinegar, which contains ethanoic acid, CH₃COOH
• Lemon juice, which contains citric acid
• Carbonated drinks, in which dissolved carbon dioxide contributes to carbonic acid equilibria
• Stomach contents, which contain hydrochloric acid

For ethanoic acid in water:

CH₃COOH(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq)

Ethanoic acid donates a proton and becomes ethanoate, CH₃COO⁻. Because the reaction is incomplete, ethanoic acid is classified as a weak acid. Its identity as an acid depends on what it does, while its strength depends on how extensively it ionizes.

What makes a substance a base?

A Brønsted-Lowry base is a species capable of accepting a proton. A base normally needs an available lone pair of electrons or a negatively charged region that can form a bond to H⁺.

Ammonia provides an important example:

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

NH₃ accepts a proton from water, so NH₃ is the base. Water donates that proton, so water acts as the acid. The reaction produces ammonium ions and hydroxide ions, making the solution basic.

This example shows why the statement “a base contains OH⁻” is incomplete. Ammonia contains no hydroxide ion, but it is still a base because it accepts H⁺. Hydroxide ions appear because ammonia removes protons from water.

Everyday base examples

Common basic substances include:

• Baking soda, containing sodium hydrogencarbonate, NaHCO₃
• Household ammonia solutions, containing NH₃ dissolved in water
• Some soaps and detergents, which produce basic solutions
• Antacids, which may contain bases such as magnesium hydroxide or calcium carbonate

The hydrogencarbonate ion in baking soda accepts a proton from an acid:

HCO₃⁻(aq) + H⁺(aq) → H₂CO₃(aq)

Carbonic acid can then decompose:

H₂CO₃(aq) → CO₂(g) + H₂O(l)

This explains the bubbling observed when baking soda reacts with vinegar. The hydrogencarbonate ion acts as a base by accepting a proton from ethanoic acid, and carbon dioxide gas is subsequently released.

Conjugate acid-base pairs

Every Brønsted-Lowry proton transfer produces two conjugate acid-base pairs. Members of a conjugate pair differ by exactly one proton.

The general reaction is:

HA + B ⇌ A⁻ + BH⁺

• HA donates H⁺ and becomes A⁻.
• HA and A⁻ form one conjugate acid-base pair.
• B accepts H⁺ and becomes BH⁺.
• BH⁺ and B form the second conjugate acid-base pair.

The IUPAC definition of a conjugate acid-base pair emphasizes this one-proton relationship.

For the reaction between ammonia and water:

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

A quick exam method is to compare the formulas. The acid has one more H and a charge one unit more positive than its conjugate base. For example, NH₄⁺ and NH₃ differ by H⁺, as do H₂CO₃ and HCO₃⁻.

For more practice identifying these relationships, use RevisionDojo's conjugate acid-base pairs explanation.

Why the same substance can act as an acid or a base

Acid-base identity can depend on the other reactant. A species that can both donate and accept a proton is called amphiprotic.

Water is the most important example. With HCl, water accepts a proton and acts as a base:

HCl + H₂O → H₃O⁺ + Cl⁻

With ammonia, water donates a proton and acts as an acid:

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

The hydrogencarbonate ion is also amphiprotic. It can accept H⁺ to form H₂CO₃ or donate H⁺ to form CO₃²⁻:

HCO₃⁻ + H⁺ → H₂CO₃

HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O

This contextual behavior is why students should identify acids and bases from a specific equation, rather than assuming that each formula always has one fixed role. The OpenStax explanation of Brønsted-Lowry acids and bases provides further examples of amphiprotic behavior.

How pH distinguishes acidic and basic solutions

The pH scale describes the hydrogen ion concentration of an aqueous solution:

pH = −log₁₀[H⁺]

IB commonly writes [H⁺], although [H₃O⁺] is a more explicit representation in water. Because pH is logarithmic, a change of one pH unit corresponds to a tenfold change in hydrogen ion concentration.

For example:

• A solution with pH 3 has [H⁺] = 1 × 10⁻³ mol dm⁻³.
• A solution with pH 4 has [H⁺] = 1 × 10⁻⁴ mol dm⁻³.
• The pH 3 solution therefore has ten times the hydrogen ion concentration of the pH 4 solution.

At 25 °C, aqueous solutions are classified as follows:

The temperature qualification matters. Neutrality always means [H₃O⁺] = [OH⁻], but neutral pH is not necessarily exactly 7 at temperatures other than 25 °C. See RevisionDojo's explanation of the pH of a neutral solution at 25 °C and the OpenStax pH and pOH reference for the quantitative relationships.

Acid or base is not the same as strong or weak

The difference between an acid and a base concerns the direction of proton transfer. Strength concerns the extent of ionization or proton transfer in a particular solvent.

• A strong acid ionizes essentially completely in water.
• A weak acid ionizes only partially and establishes an equilibrium.
• A strong base reacts essentially completely with available protons or dissociates completely to supply the basic species.
• A weak base accepts protons only partially and establishes an equilibrium.

For example, HCl is a strong acid in water:

HCl + H₂O → H₃O⁺ + Cl⁻

Ethanoic acid is weak:

CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻

The one-way arrow for HCl represents essentially complete ionization, while the equilibrium arrow for ethanoic acid represents partial ionization. RevisionDojo covers this distinction in more detail in strong vs weak acids, strong vs weak bases, and Ka and Kb for IB Chemistry.

Strength is not concentration

Strength describes the proportion that ionizes or reacts. Concentration describes the amount of solute per unit volume, usually in mol dm⁻³.

A dilute strong acid can have a higher pH than a concentrated weak acid. You therefore cannot identify strength from pH unless concentration and other relevant conditions are controlled. The same warning applies to bases.

This is one of the most common acid-base errors in IB exams. If the question asks why an acid is strong, refer to essentially complete ionization, not merely a low pH.

What happens when an acid reacts with a base?

An acid-base reaction transfers a proton from the acid to the base. In a familiar strong acid-strong base neutralization, the net ionic equation is:

H⁺(aq) + OH⁻(aq) → H₂O(l)

For example:

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

HCl supplies the acidic proton, while OH⁻ from sodium hydroxide accepts it. Na⁺ and Cl⁻ are spectator ions in the net ionic equation.

However, acid-base reactions do not always produce water. The reaction between HCl and NH₃ is also a proton transfer:

HCl + NH₃ → NH₄⁺ + Cl⁻

NH₃ accepts H⁺ directly, producing NH₄⁺. Defining neutralization only as “acid plus base gives salt and water” is therefore too narrow for Brønsted-Lowry theory.

How indicators reveal acidic and basic conditions

An acid-base indicator is generally a weak acid or weak base whose protonated and deprotonated forms have different colors. Its color therefore provides evidence about the pH range of a solution, rather than directly proving that a particular compound is intrinsically strong or weak.

Red cabbage pigment can also act as a natural indicator. The American Chemical Society red cabbage activity identifies vinegar and lemon juice as acidic examples and baking soda solution as a basic example.

Indicators should be interpreted carefully. A specific indicator changes color only across its own transition range, so it does not necessarily provide an exact pH. RevisionDojo's guide to how acid-base indicators work explains the equilibrium behind the color change.

Never identify laboratory acids or bases by tasting or touching them. Both concentrated acids and concentrated bases can be corrosive, and familiar descriptions such as “sour” or “slippery” are observations, not safe testing procedures.

How to identify an acid and a base in an IB exam

When given a chemical equation, use a systematic proton-tracking method:

1. Compare each reactant with the related product. Look for a change of one H⁺.
2. Identify the species that loses H⁺. That reactant is the acid.
3. Identify the species that gains H⁺. That reactant is the base.
4. Match the conjugate pairs. Each pair must differ by exactly one proton.
5. Check the charges. Losing H⁺ makes the charge one unit more negative; gaining H⁺ makes it one unit more positive.

Consider:

HSO₄⁻ + NH₃ ⇌ SO₄²⁻ + NH₄⁺

HSO₄⁻ becomes SO₄²⁻ after losing H⁺, so HSO₄⁻ is the acid. NH₃ becomes NH₄⁺ after gaining H⁺, so NH₃ is the base. The conjugate pairs are HSO₄⁻/SO₄²⁻ and NH₄⁺/NH₃.

For broader syllabus review without duplicating this single-concept explanation, use RevisionDojo's topic-wide IB Chemistry acids and bases exam-focused resources. Its linked notes, flashcards, and question practice can help you move from definitions to calculations and unfamiliar equations.

Common mistakes to avoid

• Defining every acid as a substance containing hydrogen. CH₄ contains hydrogen but does not normally donate H⁺ in aqueous acid-base chemistry.
• Defining every base as a substance containing OH⁻. NH₃ is a base because it accepts H⁺.
• Assigning roles without using the equation. Amphiprotic species can change roles depending on their reaction partner.
• Pairing species that differ by more than one proton. Conjugate pairs differ by exactly one H⁺.
• Confusing strong with concentrated. Strength is extent of ionization; concentration is amount per volume.
• Assuming neutral always means pH 7. That numerical value applies to neutral water at 25 °C.
• Calling any reaction that produces a salt a neutralization. The defining process is proton transfer from an acid to a base.

Conclusion

The fundamental difference between an acid and a base is the direction in which a proton moves: an acid donates H⁺, while a base accepts H⁺. The acid becomes its conjugate base, and the base becomes its conjugate acid, with each conjugate pair differing by one proton.

For IB Chemistry, apply this definition directly to equations rather than relying on pH, formula patterns, or everyday properties alone. Once the roles are secure, use pH to describe aqueous conditions and ionization extent to distinguish strong from weak species. RevisionDojo's acids and bases notes, flashcards, Questionbank, and Jojo AI can then be used to practise proton tracking and receive feedback on exam-style explanations.

Sources and referenced URLs

• IB Diploma Programme Chemistry subject brief, first assessment 2025
• IUPAC Gold Book definition of a Brønsted acid
• IUPAC Gold Book definition of a conjugate acid-base pair
• OpenStax Chemistry 2e: Brønsted-Lowry acids and bases
• OpenStax Chemistry 2e: pH and pOH
• American Chemical Society red cabbage indicator activity
• RevisionDojo IB Chemistry acids and bases resources
• RevisionDojo conjugate acid-base pairs explanation
• RevisionDojo strong vs weak acids explanation
• RevisionDojo strong vs weak bases explanation
• RevisionDojo Ka and Kb explanation
• RevisionDojo pH of a neutral solution at 25 °C
• RevisionDojo acid-base indicators explanation