Difficulty: Introductory | Prerequisites: Basic chemical bonding, molecular structure
There are three ways to define acids and bases, each broader than the last: Arrhenius (H⁺ and OH⁻ in water), Brønsted-Lowry (proton donors and acceptors, any solvent), and Lewis (electron-pair acceptors and donors). Every Arrhenius acid is a Brønsted acid, and every Brønsted acid is a Lewis acid, but not the other way round. Some species, like water, are amphoteric, meaning they can act as either an acid or a base depending on their environment.
Arrhenius acid
A substance that dissociates in aqueous solution to produce an excess of H⁺ ions. In simple terms, it is anything that makes water more acidic by releasing hydrogen ions.
Arrhenius base
A substance that dissociates in aqueous solution to produce an excess of OH⁻ ions. Think of it as anything that makes water more basic by releasing hydroxide ions.
Brønsted-Lowry acid
A species that donates a proton (H⁺) to another species. This definition works in any solvent, not just water.
Brønsted-Lowry base
A species that accepts a proton (H⁺) from another species. In simple terms, it is the molecule or ion that "catches" the proton the acid gives up.
Lewis acid
An electron-pair acceptor. Think of it as a species with an empty orbital that can receive a pair of electrons from something else.
Lewis base
An electron-pair donor. It uses a lone pair to form a new sigma bond (also called a coordinate covalent bond or dative bond) with a Lewis acid.
Conjugate acid-base pair
Two species that differ by exactly one proton. When an acid donates H⁺, the product is its conjugate base; when a base accepts H⁺, the product is its conjugate acid.
Amphoteric
A species that can act as an acid in a basic environment and as a base in an acidic environment. Water is the classic example.
Amphiprotic
A subset of amphoteric: a species that can either donate or accept a proton. All amphiprotic species are amphoteric, but not all amphoteric species are amphiprotic (metal oxides are amphoteric but have no protons to donate).
Oxyacid
An acid formed from an oxyanion (an anion containing oxygen). Naming depends on the number of oxygens: anions ending in -ite give -ous acids, and anions ending in -ate give -ic acids.
Zwitterion
A molecule that carries both a positive and a negative charge at different sites. Zwitterions are amphoteric because each charged group can interact with acids or bases.
Acids dissociate in water to produce excess H⁺; bases dissociate to produce excess OH⁻
Only applies to aqueous solutions
Naming rule: acids must have H at the beginning of their formula, bases must have OH at the end
Water itself is not an Arrhenius acid because it does not produce an excess of H⁺
Acid = proton (H⁺) donor; base = proton acceptor
Works in any solvent, not limited to water
Introduces conjugate acid-base pairs: every acid-base reaction produces a conjugate acid and a conjugate base
Water is a Brønsted acid (it can donate a proton) and also a Brønsted base (it can accept one)
F⁻ is a Brønsted base but not an Arrhenius base, because it accepts a proton rather than dissociating to release OH⁻
Acid = electron-pair acceptor; base = electron-pair donor
The donated pair forms a new coordinate covalent bond (sigma bond)
Captures species the other models miss: BF₃ is a Lewis acid (it accepts an electron pair) but it is neither a Brønsted nor an Arrhenius acid because it has no proton to donate and does not produce H⁺ in water
Lewis acids are common catalysts in organic chemistry
All Arrhenius acids are Brønsted acids, and all Brønsted acids are Lewis acids. The reverse is not true. For example, NH₃ is a Brønsted base but not an Arrhenius base (it does not dissociate to release OH⁻).
Can behave as an acid or a base depending on the environment
Water is the textbook example: it donates a proton to strong bases and accepts one from strong acids
Amphiprotic is a narrower term: the species can both lose and gain a proton. All amphiprotic species are amphoteric, but metal oxides and hydroxides are amphoteric without being amphiprotic (they have no protons to give off)
Partially dissociated conjugate bases of polyvalent acids are usually amphoteric (e.g. HSO₄⁻ can either gain a proton or lose its remaining one)
Zwitterions are also amphoteric
Anion ends in -ide: the acid gets the prefix hydro- and the suffix -ic (e.g. fluoride becomes hydrofluoric acid)
Anion is an oxyanion ending in -ite (one or two oxygens): the acid ends in -ous (e.g. sulfite becomes sulfurous acid)
Anion is an oxyanion ending in -ate (three or four oxygens): the acid ends in -ic (e.g. sulfate becomes sulfuric acid)
Acids formed from oxyanions are called oxyacids
The Lewis acid-base framework is the one that matters most in organic chemistry: Lewis acids like AlCl₃ and BF₃ catalyse Friedel-Crafts reactions and other electrophilic processes. The bicarbonate buffer system in blood relies on amphiprotic behaviour to regulate pH, which is why hospitals monitor blood CO₂ levels.
Students often assume the three acid-base definitions are interchangeable. They are nested: Arrhenius ⊂ Brønsted-Lowry ⊂ Lewis. A Lewis acid (e.g. BF₃) may have nothing to do with protons.
"Amphoteric" and "amphiprotic" are frequently confused. Amphoteric is the broader term (acts as acid or base). Amphiprotic specifically means the species can donate or accept a proton. Metal oxides are amphoteric but not amphiprotic.
Students sometimes think water cannot be an acid. Under the Brønsted-Lowry definition it can donate a proton, so it qualifies. Under Arrhenius it does not, because it does not produce an excess of H⁺.
The nomenclature suffixes -ous/-ic are often mixed up. The rule: -ite anions (fewer oxygens) give -ous acids, -ate anions (more oxygens) give -ic acids. A mnemonic: "-ite" is a smaller word, "-ous" has fewer letters than "-ic acid" when you include the full name.
⚠️ Expect questions that give you a species (e.g. NH₃, BF₃, F⁻) and ask which acid-base model(s) it fits. Know the hierarchy cold.
⚠️ Conjugate acid-base pair identification is heavily tested. Be ready to identify the conjugate base of any given acid, and vice versa.
⚠️ Nomenclature questions ("name the acid formed from the perchlorate ion") appear on nearly every general chemistry exam. Memorise the -ite/-ous and -ate/-ic pattern and the hydro-/-ic pattern for binary acids.
⚠️ Amphoteric species questions often ask you to write two reactions for the same species: one where it acts as an acid, one where it acts as a base. Water and HSO₄⁻ are the most common examples.
True or false: All Brønsted-Lowry acids are also Arrhenius acids. (Answer: False. Brønsted-Lowry is broader.)
Fill in the blank: A Lewis base donates a(n) ______ to a Lewis acid. (Answer: electron pair)
True or false: HSO₄⁻ is amphiprotic. (Answer: True. It can donate or accept a proton.)
Fill in the blank: An acid formed from the chloride ion (Cl⁻) is called ______ acid. (Answer: hydrochloric)
True or false: BF₃ is a Brønsted-Lowry acid. (Answer: False. It is a Lewis acid only.)
Q: Classify NH₃ under each of the three acid-base definitions. Is it an acid, a base, both, or neither in each framework?
A: Arrhenius: neither (it does not dissociate to produce OH⁻ directly). Brønsted-Lowry: base (it accepts a proton). Lewis: base (it donates its lone pair).
Q: Write the conjugate acid of F⁻ and the conjugate base of H₂O.
A: Conjugate acid of F⁻ is HF (F⁻ gains a proton). Conjugate base of H₂O is OH⁻ (H₂O loses a proton).
Q: Name the acid formed from the sulfite ion (SO₃²⁻).
A: Sulfurous acid (H₂SO₃). The anion ends in -ite, so the acid ends in -ous.
Q: Explain why BF₃ is a Lewis acid but not a Brønsted-Lowry acid.
A: BF₃ has an empty p-orbital and accepts an electron pair, which makes it a Lewis acid. It has no proton (H⁺) to donate, so it cannot be a Brønsted-Lowry acid.
Q: Water can act as both an acid and a base. What term describes this property, and give the Brønsted-Lowry reaction in which water acts as an acid.
A: Water is amphoteric (and specifically amphiprotic). As an acid: H₂O + NH₃ → OH⁻ + NH₄⁺ (water donates a proton to ammonia).
This material connects directly to pH, pOH, and the autoionization of water (covered in Part 2 of these notes), because you need to know what makes something acidic or basic before you can quantify how acidic or basic it is. The Lewis model reappears heavily in organic chemistry, where electrophiles (Lewis acids) and nucleophiles (Lewis bases) drive nearly every reaction mechanism. Amphoteric behaviour ties into buffer chemistry and the bicarbonate buffer system in physiology.
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