Unit 8: Acids and Bases

Arrhenius acid

A substance that produces hydronium (H3O+) in water (often written as producing H+).

Arrhenius base

A substance that produces hydroxide (OH−) in water.

Brønsted–Lowry acid

A proton (H+) donor.

Brønsted–Lowry base

A proton (H+) acceptor.

Lewis acid

An electron-pair acceptor (broader than Brønsted–Lowry; may involve no protons).

Lewis base

An electron-pair donor (e.g., NH3 donates a lone pair to accept a proton).

Conjugate base

The species formed when an acid donates a proton (acid minus H+).

Conjugate acid

The species formed when a base accepts a proton (base plus H+).

Conjugate acid–base pair

Two species that differ by exactly one proton (H+).

Amphiprotic species

A substance that can act as either an acid or a base depending on what it reacts with (e.g., H2O).

Acid strength (Unit 8 meaning)

How completely an acid reacts with water (equilibrium position), not how concentrated it is.

Strong acid

Reacts essentially completely with water to form H3O+; in calculations, [H3O+] is determined by stoichiometry.

Weak acid

Only partially ionizes in water; requires equilibrium (Ka/ICE table) to find [H3O+].

Strong base

Dissociates essentially completely in water to produce OH− (e.g., NaOH); [OH−] is determined by stoichiometry.

Weak base

Partially reacts with water by accepting a proton, producing OH− (requires Kb/ICE table).

Common strong acids (AP list)

HCl, HBr, HI, HNO3, HClO4, H2SO4 (first dissociation treated as strong).

Common strong bases (AP list)

LiOH, NaOH, KOH, Ba(OH)2, Sr(OH)2.

Autoionization of water

2H2O ⇌ H3O+ + OH− (pure water ionizes slightly).

Kw (ion-product constant of water)

Kw = [H3O+][OH−] = 1.0×10−14 at 25°C; pKw = 14.00 at 25°C.

pH

pH = −log[H3O+] (base-10 log; [H+] is shorthand for [H3O+]).

pOH

pOH = −log[OH−] (base-10 log).

Relationship between pH and pOH (25°C)

pH + pOH = 14.00 at 25°C (because pH + pOH = pKw).

Ka (acid dissociation constant)

Equilibrium constant for acid ionization; larger Ka means stronger acid.

Kb (base dissociation constant)

Equilibrium constant for base reacting with water to form OH−; larger Kb means stronger base.

pKa

pKa = −log Ka (smaller pKa means stronger acid).

pKb

pKb = −log Kb (smaller pKb means stronger base).

Conjugate relationship: KaKb = Kw

For a conjugate acid–base pair, Ka(HA)·Kb(A−) = Kw; at 25°C, pKa + pKb = 14.00.

ICE table

A setup (Initial, Change, Equilibrium) used to solve equilibrium concentrations for weak acids/bases.

Weak-acid equilibrium (HA)

HA + H2O ⇌ H3O+ + A−; Ka = ([H3O+][A−])/[HA].

Weak-base equilibrium (B)

B + H2O ⇌ BH+ + OH−; Kb = ([BH+][OH−])/[B].

5% approximation check

If x/C × 100% < 5%, you may approximate C − x ≈ C in Ka or Kb ICE-table calculations; otherwise solve the quadratic.

Percent ionization (percent dissociation)

% ionization = ([H3O+]eq / [HA]initial)×100%; for weak acids, it increases as the solution is diluted.

Neutralization (net ionic)

H3O+ + OH− → 2H2O (key reaction when strong acid and strong base are mixed).

Strong acid–strong base mixing workflow

Convert to moles → do neutralization stoichiometry → find excess acid/base → divide by total volume → convert to pH/pOH.

Polyprotic acid

An acid that can donate more than one proton (e.g., H2SO4, H3PO4), typically dissociating stepwise.

Stepwise dissociation constants (Ka1, Ka2, …)

Each proton loss has its own Ka; usually Ka1 > Ka2 > Ka3 because losing later protons is less favorable.

Amphiprotic intermediate ion

An intermediate from a polyprotic acid (e.g., HA−) that can act as an acid or a base in water.

Conjugate-base stability principle

Stronger acids have more stable conjugate bases; stability of the conjugate base drives acidity trends.

Resonance stabilization (acidity)

Delocalization of negative charge in the conjugate base stabilizes it and increases acid strength (e.g., carboxylic acids > alcohols).

Inductive effect (acidity)

Electron-withdrawing groups stabilize negative charge through sigma bonds, strengthening the acid (e.g., CF3COOH > CH3COOH).

Down-group H–X acidity trend

For binary acids H–X down a group, acidity increases as X gets larger because H–X bonds weaken and negative charge is better stabilized (e.g., HI > HF).

Oxyacid oxygen-number trend

With the same central atom, more oxygens usually means a stronger oxyacid (greater resonance/inductive stabilization of the conjugate base), e.g., HClO3 > HClO2 > HClO.

Oxyacid central-atom electronegativity trend

With the same number of oxygens, a more electronegative central atom tends to make a stronger oxyacid (stabilizes the conjugate base via induction).

Common ion effect (hydroxide solubility)

Adding OH− shifts M(OH)n(s) ⇌ Mn+ + nOH− left (less dissolves); adding H+ removes OH− as H2O, pulling dissolution right (more dissolves).

Salt hydrolysis

Reaction of dissolved ions with water that makes a solution acidic or basic; conjugate acids of weak bases (e.g., NH4+) produce H3O+, and conjugate bases of weak acids (e.g., F−, C2H3O2−) produce OH−.

Equivalence point vs endpoint

Equivalence point: stoichiometric point where moles acid and base react exactly; endpoint: indicator color change observed experimentally (should be close to equivalence).

Acid–base indicator

A weak acid/base whose conjugate forms have different colors; color transition occurs near its pKa (often ~pKa ± 1), so choose an indicator whose transition range lies in the steep part of the titration curve.

Buffer

A mixture (typically weak acid/conjugate base or weak base/conjugate acid) that resists pH change by consuming added H3O+ or OH− via conjugate reactions.

Henderson–Hasselbalch equation

For a weak-acid buffer: pH = pKa + log([A−]/[HA]); after adding strong acid/base, do stoichiometry first, then apply this to the new ratio.

Half-equivalence point

In a weak acid–strong base titration, the point where [HA] = [A−], so pH = pKa (lets you read pKa from the curve).