Acids and bases — Arrhenius, Brønsted-Lowry, Lewis definitions

The concept of acids and bases has evolved through three major definitions, each expanding our understanding beyond the previous one. These aren't competing theories—they're nested frameworks where each definition is more general than the last.

Why multiple definitions?

Think of it like defining "vehicle": first you define "car" (Arrhenius), then "any land transport" (Brønsted-Lowry), then "anything that moves people or goods" (Lewis). Each is valid in its domain.

1. Arrhenius Definition (1884)

An Arrhenius base is a substance that produces $\text{OH}^-$ ions when dissolved in water.

Limitation: Only applies to aqueous solutions. Can't explain ammonia's basicity in non-aqueous solvents or $\text{NH}_3$ acting as a base despite having no $\text{OH}^-$ to donate.

Derivation of neutralization

When an Arrhenius acid meets an Arrhenius base in water:

$\text{H}^+ \text{ (aq)} + \text{OH}^- \text{ (aq)} \rightarrow \text{H}_2\text{O (l)}$

Why does this happen? The $\text{H}^+$ and $\text{OH}^-$ ions have extremely high attraction (strong electrostatic force) and form the very stable water molecule ($\Delta H = -55.8 \text{ kJ/mol}$). This is why acid-base neutralization is always exothermic.

For example, hydrochloric acid and sodium hydroxide:

$\text{HCl (aq) + NaOH (aq) → NaCl (aq) + H}_2\text{O (l)}$

The net ionic equation strips away spectator ions ($\text{Na}^+$, $\text{Cl}^-$):

$\text{H}^+ \text{ (aq)} + \text{OH}^- \text{ (aq)} \rightarrow \text{H}_2\text{O (l)}$

Why this step? Spectator ions don't participate in the actual reaction—they're already dissociated and remain dissociated. The essence of Arrhenius neutralization is always $\text{H}^+$ meeting $\text{OH}^-$.

$\text{H}_2\text{SO}_4\text{ (aq)} \to 2\text{H}^+\text{ (aq)} + \text{SO}_4^{2-}\text{ (aq)}$

Why this step? Sulfuric acid is diprotic—each molecule can donate 2 protons. The first proton dissociates completely (strong acid behavior), the second partially.

If we neutralize with $\text{NaOH}$:

$\text{H}_2\text{SO}_4\text{ (aq) + 2NaOH (aq) } \rightarrow \text{ Na}_2\text{SO}_4\text{ (aq) + 2H}_2\text{O (l)}$

Why this step? We need 2 moles of $\text{OH}^-$ to neutralize 2 moles of $\text{H}^+$. Stoichiometry is 1:2.

$NH_3 \text{ (aq)} + H_2O \text{ (l)} \rightleftharpoons NH_4^+ \text{ (aq)} + OH^- \text{ (aq)}$

Why this step? Ammonia abstracts a proton from water, generating $\text{OH}^-$ ions indirectly. So it's technically an Arrhenius base because it produces $\text{OH}^-$ in water—but this explanation feels forced. This limitation motivated Brønsted-Lowry.

2. Brønsted-Lowry Definition (1923)

A Brønsted-Lowry base is a proton (H⁺) acceptor.

Key insight: Focuses on proton transfer, not ion production in water. Works in any solvent, even gas phase.

Conjugate acid-base pairs

Every Brønsted-Lowry acid-base reaction involves two conjugate pairs. When an acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid.

Why? Because the reaction is reversible. If $\text{HA}$ donates a proton to become $\text{A}^-$, then $\text{A}^-$ can accept a proton to become $\text{HA}$ again.

General form:

$\text{HA + B} \rightleftharpoons \text{A}^- + \text{HB}^+$

• $\text{HA}$ (acid) and $\text{A}^-$ (conjugate base) differ by one $\text{H}^+$
• $\text{B}$ (base) and $\text{HB}^+$ (conjugate acid) differ by one $\text{H}^+$

Identify the pairs:

• $\text{H}_2\text{O}$ (acid) donates $\text{H}^+$ → $\text{OH}^-$ (conjugate base)
• $\text{NH}_3$ (base) accepts $\text{H}^+$ → $\text{NH}_4^+$ (conjugate acid)

Why this step? Water acts as an acid here (amphoteric behavior). Ammonia's lone pair on nitrogen accepts the proton. This explains ammonia's basicity without needing $\text{OH}^-$ production as the definition.

• $\text{HCl}$ (acid) → $\text{Cl}^{-}$ (conjugate base)
• $\text{H}_2\text{O}$ (base) → $\text{H}_3\text{O}^+$ (conjugate acid)

Why this step? Water accepts the proton from $\text{HCl}$, forming hydronium. This is more accurate than saying $\text{H}^+$ exists freely—protons are always solvated.

Strong vs weak: $\text{HCl}$ is a strong acid because the forward reaction goes to completion. Weak acids establish equilibrium. The strength of an acid is inversely related to its conjugate base strength: strong acid → weak conjugate base.

Amphoteric species

Water is the classic example:

• As acid: $\text{H}_2\text{O} + \text{NH}_3 \rightarrow \text{OH}^- + \text{NH}_4^+$
• As base: $H_2O + HCl \rightarrow H_3O^+ + Cl^-$

Why? Water has both a proton to donate ($\text{H}$ from $\text{H-O-H}$) and lone pairs on oxygen to accept a proton.

Acting as acid: $\ce{HCO3- + OH- -> CO3^2- + H2O}$

Acting as base: $\ce{HCO3- + H+ -> H2CO3}$

Why this step? Bicarbonate has a removable proton (acid behavior) and oxygen lone pairs that can accept a proton (base behavior). This makes it an excellent buffer.

3. Lewis Definition (1923)

A Lewis base is an electron-pair donor.

Key insight: No protons required. Includes all Brønsted-Lowry acids/bases plus reactions involving only electron-pair transfers (like metal-ligand complexation).

Why this generalization works

In Brønsted-Lowry, the base donates electrons (as a lone pair) to form a bond with $\ce{H+}$. Lewis simply removes the requirement that it must be a proton—any electrophile can be the acceptor.

$\ce{A + :B -> A-B}$

• $\ce{A}$ = Lewis acid (electron-deficient)
• $\ce$ = Lewis base (has lone pair)
• $\ce{A-B}$ = aduct (coordinate covalent bond)

• $\ce{BF3}$ is a Lewis acid: boron has only6 valence electrons (incomplete octet), accepting pair from nitrogen.
• $\ce{NH3}$ is a Lewis base: nitrogen has a lone pair to donate.

Why this step? The nitrogen lone pair forms a dative bond (coordinate covalent) with the empty p-orbital on boron. No protons involved, but it's still an acid-base reaction by Lewis.

• $\ce{Ag+}$ is a Lewis acid (empty orbitals, positive charge attracts electrons)
• $\ce{NH3}$ is a Lewis base (lone pair donor)

Why this step? Silver ion coordinates with ammonia through donation of nitrogen's lone pairs into silver's empty orbitals. This forms a coordination complex. All metal complexation is Lewis acid-base chemistry.

Real-world: This reaction is used to dissolve silver chloride in qualitative analysis.

Through the Lewis lens:

• $\ce{H+}$ (a bare proton) is the ultimate Lewis acid—it has no electrons, only wants them.
• $\ce{H2O}$ is a Lewis base (oxygen lone pairs).

Why this step? This shows that Brønsted-Lowry is a subset of Lewis. Every proton transfer is an electron-pair donation to $\ce{H+}$.

Comparison and when to use each

Use Arrhenius for simple aqueous acid-base titrations and reactions.

Use Brønsted-Lowry for buffer chemistry, conjugate pairs, proton transfer mechanisms.

Use Lewis for coordination chemistry, catalysis, and reactions with no protons (e.g., $\ce{CO2}$ as acid: $\ce{CO2 + O^2- -> CO3^2-}$).

The fix: Arrhenius is perfectly correct within its domain (aqueous solutions). It's not wrong—it's just limited. All Arrhenius acids/bases are Brønsted-Lowry acids/bases, but not vice versa. Similarly, all Brønsted-Lowry acids/bases are Lewis acids/bases. They're nested, not contradictory.

The fix: All Lewis acids are electrophiles, but "Lewis acid" specifically refers to the formation of a coordinate covalent bond (aduct). In contrast, electrophiles in organic chemistry often form regular covalent bonds through electron transfer. It's a terminology distinction—Lewis is broader and includes coordination.

The fix: In $\ce{HCl + H2O -> H3O+ + Cl-}$, water is the base and hydronium is its conjugate acid. In $\ce{NH3 + H2O -> NH4+ + OH-}$, water is the acid and hydroxide is its conjugate base. Water's amphoteric nature is central to aqueous equilibrium.

First way (Arrhenius): A helper is someone who brings snacks ($\ce{H+}$) or juice boxes ($\ce{OH-}$) to lunch. Simple, but what about someone who helps with homework? They don't bring food.

Second way (Brønsted-Lowry): A helper is anyone who shares something (proton donor) or receives something shared (proton acceptor). Now we can include the homework helper who accepts your questions and shares answers. Works anywhere, not just at lunch.

Third way (Lewis): A helper is anyone who gives anything (electrons) or receives anything. This includes the homework helper, the snack sharer, and the kid who holds your backpack while you tie your shoes. Super broad—covers everything.

Acids and bases are like this. Scientists kept expanding the definition as they found more examples that didn't fit the old rules. Each definition is still useful depending on what you're studying!

Or: "All Birds Learn" (Arrhenius → Brønsted → Lewis, each broader than the last)

Connections

• Le Chatelier's Principle — explains why adding $\ce{H+}$ or $\ce{OH-}$ shifts conjugate pair equilibria
• pH and pOH — quantifies Arrhenius $\ce{H+}$ and $\ce{OH-}$ concentrations
• Buffer solutions — rely on Brønsted-Lowry conjugate pairs
• Acid-base equilibrium constants — $K_a$, $K_b$, $K_w$ derive from these definitions
• Coordination compounds — entirely based on Lewis acid-base theory
• Hydrolysis of salts — uses conjugate acid-base strength relationships
• Amphoteric oxides — substances that react with both acids and bases

#flashcards/chemistry

What is an Arrhenius acid? :: A substance that produces $\ce{H+}$ (or $\ce{H3O+}$) ions when dissolved in water.

What is the relationship between acid strength and conjugate base strength? :: Inversely related. A strong acid has a weak conjugate base; a weak acid has a strong conjugate base.

Concept Map

Hinglish (regional understanding)