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  1. AP Chemistry

  3. Introduction to Acids and Bases

AP CHEMISTRY • ACIDS AND BASES

Introduction to Acids and Bases

Understanding proton transfer and electron-pair acceptance as the foundation of chemical reactivity in aqueous systems.

SECTION 1

Historical Context & Motivation

The classification of substances as acids or bases is one of the oldest organizing schemes in chemistry, rooted in practical observations about taste, corrosivity, and the ability of certain solutions to neutralize one another. Early alchemists noted that acidic substances turned blue litmus red and dissolved metals, while alkaline solutions felt slippery and reversed litmus color changes. These empirical distinctions persisted for centuries before chemists developed molecular-level explanations for why these behaviors occur.

1661
Robert Boyle's Empirical Definitions
Boyle catalogued observable properties of acids (sour taste, corrosive, turn litmus red) and bases (slippery, turn litmus blue), establishing the first systematic classification.
1884
Arrhenius Theory
Svante Arrhenius proposed that acids produce H⁺ ions and bases produce OH⁻ ions upon dissolution in water, earning him the 1903 Nobel Prize in Chemistry.
1923
Brønsted–Lowry Theory
Johannes Brønsted and Thomas Lowry independently defined acids as proton donors and bases as proton acceptors, extending acid–base chemistry beyond aqueous solutions.
1923
Lewis Theory
Gilbert N. Lewis broadened the definition further: acids accept electron pairs, bases donate them. This framework encompasses reactions with no protons at all, such as BF₃ + NH₃.

The central question driving these successive refinements was: What is the fundamental chemical event that makes a substance acidic or basic? Each model offered a wider lens—from ions in water (Arrhenius) to proton transfer in any solvent (Brønsted–Lowry) to electron-pair transactions in all phases (Lewis). The AP Chemistry curriculum emphasizes the Brønsted–Lowry framework for equilibrium calculations while invoking Lewis theory for coordination chemistry and reaction mechanisms.

SECTION 2

Core Principles & Definitions

Three theoretical frameworks coexist in modern chemistry, each useful in different contexts. The AP exam expects you to move fluently among them, recognizing when to apply each definition and understanding their nested relationship: every Arrhenius acid is also a Brønsted–Lowry acid, and every Brønsted–Lowry acid is also a Lewis acid, but the reverse is not true.

1

Arrhenius Model

An acid dissociates in water to release H⁺ (or H₃O⁺); a base dissociates to release OH⁻. Applicable only in aqueous solutions. Example: HCl → H⁺ + Cl⁻; NaOH → Na⁺ + OH⁻.
2

Brønsted–Lowry Model

An acid is a proton (H⁺) donor; a base is a proton acceptor. Every acid–base reaction involves a conjugate acid–base pair. Works in any solvent or gas phase.
3

Lewis Model

An acid is an electron-pair acceptor; a base is an electron-pair donor. This is the most general framework and includes metal-ion coordination, BF₃ + NH₃, and more.
4

Conjugate Pairs

When an acid donates a proton it becomes its conjugate base; when a base accepts a proton it becomes its conjugate acid. A strong acid has a very weak conjugate base.
5

Amphiprotic Species

Substances like water (H₂O) and the hydrogen carbonate ion (HCO₃⁻) can act as either an acid or a base depending on the reaction partner.
✦ KEY TAKEAWAY
KEY TAKEAWAY
SECTION 3

Visual Explanation — Proton Transfer in Action

Brønsted–Lowry Proton Transfer: HCl + H₂OREACTANTSHClAcid (proton donor)H₂OBase (proton acceptor)H on HCl is the transferring protonH⁺ transferPRODUCTSCl⁻Conjugate baseH₃O⁺Conjugate acidConjugate Pair RelationshipsHClacid₁pair 1Cl⁻conj. base₁H₂Obase₂pair 2H₃O⁺conj. acid₂
In this Brønsted–Lowry reaction, HCl donates a proton (H⁺) to water. Two conjugate pairs form: HCl/Cl⁻ (pair 1) and H₂O/H₃O⁺ (pair 2). A strong acid like HCl reacts essentially to completion, producing a very weak conjugate base (Cl⁻).

The diagram above illustrates the essential Brønsted–Lowry mechanism: a proton donor (HCl) transfers H⁺ to a proton acceptor (H₂O). The proton never exists freely in solution; it is immediately captured by water's lone pair to form the hydronium ion (H₃O⁺). Notice that each reactant has a partner on the product side that differs by exactly one proton—these are conjugate pairs. Recognizing conjugate pairs is critical for writing equilibrium expressions and predicting the direction of proton-transfer equilibria.

SECTION 4

Mathematical Framework — pH, pOH, and Kw

Quantifying the acidity or basicity of a solution requires a logarithmic scale because hydronium concentrations span many orders of magnitude—from roughly 10 M in concentrated acid to 10⁻¹⁵ M in concentrated base. The pH scale compresses this enormous range into a manageable 0–14 span for typical aqueous systems at 25 °C.

pH DEFINITION
pH = −log[H₃O⁺]
where [H₃O⁺] is the molar concentration of hydronium ion. A low pH (< 7) indicates an acidic solution; a high pH (> 7) indicates a basic solution.
pOH DEFINITION
pOH = −log[OH⁻]
Analogous to pH but measuring hydroxide concentration. At 25 °C, pH + pOH = 14.00.
AUTOIONIZATION OF WATER
Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ (at 25 °C)
Kw is the ion-product constant for water. In pure water at 25 °C, [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ M, giving pH = 7.00. Kw increases with temperature, so neutral pH drops below 7 at elevated temperatures.
Ka AND Kb RELATIONSHIP
Ka × Kb = Kw
For any conjugate acid–base pair, the product of the acid dissociation constant (Ka) and the base dissociation constant (Kb) equals Kw. A stronger acid (larger Ka) necessarily has a weaker conjugate base (smaller Kb).
AP Exam Tip
SECTION 5

Strong vs. Weak Acids and Bases

The distinction between strong and weak acids (or bases) is one of the most consequential classifications on the AP exam. A strong acid ionizes completely in dilute aqueous solution—there is no equilibrium to consider, and Ka is so large it is not tabulated. A weak acid only partially ionizes, establishing a dynamic equilibrium described by Ka. You must memorize the six common strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄) and the strong bases (Group 1 hydroxides and Ca(OH)₂, Sr(OH)₂, Ba(OH)₂).

Strong Acid vs. Weak Acid DissociationSTRONG ACID (HCl)Complete ionization → no equilibriumHCl → H₃O⁺ + Cl⁻Single arrow: goes to completionParticle view (0.1 M solution):H₃O⁺Cl⁻H₃O⁺Cl⁻H₃O⁺Cl⁻No HCl molecules remainWEAK ACID (CH₃COOH)Partial ionization → equilibriumCH₃COOH ⇌ H₃O⁺ + CH₃COO⁻Double arrow: equilibrium existsParticle view (0.1 M solution):HAHAHAHAHAH₃O⁺A⁻Most HA molecules intact
Left: a strong acid (HCl) ionizes completely—no undissociated molecules remain. Right: a weak acid (acetic acid) reaches equilibrium with most molecules still intact. The particle-level representations highlight that concentration of H₃O⁺ is much lower for the weak acid at the same initial concentration.
Comparison of strong and weak acids/bases
PropertyStrong Acid/BaseWeak Acid/Base
Ionization in waterComplete (→)Partial (⇌)
Ka or KbVery large (not tabulated)Small (10⁻² to 10⁻¹⁴)
[H₃O⁺] from 0.10 M soln0.10 M (equals initial conc.)≪ 0.10 M (requires ICE table)
Conjugate strengthConjugate is negligibly weakConjugate has measurable strength
Common examples (acids)HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄CH₃COOH, HF, HCN, H₂CO₃
SECTION 6

Worked Example — pH of a Weak Acid

Calculate the pH of a 0.25 M acetic acid (CH₃COOH) solution. Ka = 1.8 × 10⁻⁵.

Step 1 — Write the equilibrium expression

CH₃COOH(aq) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq). The equilibrium expression is Ka = [H₃O⁺][CH₃COO⁻] / [CH₃COOH] = 1.8 × 10⁻⁵.

Step 2 — Set up the ICE table

Let x = [H₃O⁺] at equilibrium. Initial: [CH₃COOH] = 0.25, [H₃O⁺] = 0, [CH₃COO⁻] = 0. Change: −x, +x, +x. Equilibrium: (0.25 − x), x, x.

Step 3 — Apply the 5% approximation

Check if x ≪ 0.25 by assuming 0.25 − x ≈ 0.25. Then Ka = x² / 0.25 = 1.8 × 10⁻⁵. Solving: x² = 4.5 × 10⁻⁶.
x = 2.12 × 10⁻³ M

Step 4 — Validate the approximation

(2.12 × 10⁻³ / 0.25) × 100% = 0.85%, which is well under 5%. The approximation is valid.

Step 5 — Calculate pH

pH = −log(2.12 × 10⁻³) = −(−2.674) = 2.67. This is consistent with a weak acid: the pH is below 7 but higher than it would be for a strong acid of the same concentration (which would give pH = 0.60).
pH = 2.67
SECTION 7

Strengths & Limitations of Each Model

Comparison of acid–base models
ModelStrengthsLimitations
ArrheniusSimple, intuitive; directly explains neutralization producing water; easy pH calculations for strong acids/bases in water.Limited to aqueous solutions; cannot explain why NH₃ is basic (no OH⁻ in formula); cannot address gas-phase or non-aqueous reactions.
Brønsted–LowryWorks in any solvent; explains conjugate pairs; handles amphiprotic species; directly ties to Ka/Kb equilibria tested on AP exam.Still requires a transferable proton; cannot classify Lewis acids like BF₃ or metal cations that have no protons to donate.
LewisMost general; explains coordination bonds, electrophilic reactions, and why metal cations make solutions acidic; unifies organic and inorganic chemistry.Does not inherently predict Ka values or pH; can be overly broad—many reactions could technically be classified as Lewis acid–base.
✦ KEY TAKEAWAY
KEY TAKEAWAY
SECTION 8

Connection to Advanced Topics

The fundamentals of acid–base theory laid out here form the scaffolding for several advanced AP Chemistry topics. Buffer solutions rely on conjugate acid–base pairs to resist pH changes; titration curves map pH against volume of added titrant, revealing equivalence points where moles of acid equal moles of base. Polyprotic acids like H₂SO₄ and H₃PO₄ undergo stepwise dissociation, each step with its own Ka value. Lewis acid–base theory underpins coordination chemistry, where metal ions act as Lewis acids accepting electron pairs from ligands.

Introductory concepts and their advanced extensions
This Lesson CoversAdvanced Extension
pH = −log[H₃O⁺] for monoprotic acidsHenderson–Hasselbalch equation for buffer pH: pH = pKa + log([A⁻]/[HA])
Ka expression and ICE table for one equilibriumPolyprotic acid equilibria with Ka₁ ≫ Ka₂ ≫ Ka₃; stepwise ICE tables
Strong vs. weak classificationTitration curves: strong/strong, strong/weak, weak/weak; indicator selection
Lewis acid = electron-pair acceptorMetal aqua-ion hydrolysis: [Fe(H₂O)₆]³⁺ donates H⁺ making solutions acidic

As you progress through the Acids and Bases unit, every new topic will reference the concepts from this lesson: conjugate pairs, Ka/Kb relationships, the autoionization of water, and the distinction between strong and weak. Mastering these fundamentals now will make buffer and titration problems significantly more approachable.

SECTION 9

Practice Problems

PROBLEM 1 — CONCEPTUAL
In the reaction NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq), which species is the Brønsted–Lowry acid?
PROBLEM 2 — BASIC CALCULATION
What is the pH of a 0.0050 M HNO₃ solution at 25 °C?
PROBLEM 3 — INTERMEDIATE
A 0.10 M solution of a monoprotic weak acid HA has a pH of 3.00 at 25 °C. What is the Ka of HA?
PROBLEM 4 — APPLIED
Formic acid (HCOOH) has Ka = 1.8 × 10⁻⁴. A student prepares a 0.40 M HCOOH solution. (a) Write the equilibrium expression for the dissociation of formic acid in water. (b) Calculate the equilibrium concentration of H₃O⁺ and the pH of the solution. (c) Calculate the percent ionization of formic acid. (d) The student dilutes the solution to 0.040 M. Without performing a full calculation, predict whether the percent ionization will increase, decrease, or stay the same. Justify your answer.
PROBLEM 5 — CRITICAL THINKING
A student measures the pH of four 0.10 M solutions at 25 °C and records the following data: Solution A: pH = 1.00 Solution B: pH = 2.87 Solution C: pH = 4.74 Solution D: pH = 13.00 (a) Classify each solution as a strong acid, weak acid, weak base, or strong base. Justify each classification. (b) Calculate Ka for Solution C. (c) Identify the conjugate base of Solution C's acid and calculate its Kb. (d) If equal volumes of Solution A and Solution D are mixed, determine the pH of the resulting solution. Assume volumes are additive.
SUMMARY

Lesson Summary

Varsity Tutors • AP Chemistry • Introduction to Acids and Bases