Unit 8 Comprehensive Guide: Acid-Base Equilibria and Titrations

Arrhenius Acid

Produces H+ (or H3O+) in water.

Arrhenius Base

Produces OH- in water.

Brønsted-Lowry Acid

A proton (H+) donor.

Brønsted-Lowry Base

A proton (H+) acceptor.

Conjugate Acid-Base Pairs

Pairs that differ by exactly one H+.

Autoionization of Water

Process where water acts as both acid and base: 2H2O ⇌ H3O+ + OH-.

Kw

The equilibrium constant for water's autoionization: Kw = [H3O+][OH-].

pH

The negative logarithm of the hydrogen ion concentration: pH = -log[H+].

pOH

The negative logarithm of the hydroxide ion concentration: pOH = -log[OH-].

pH + pOH

Equals 14.00 at 25°C.

Strong Acids

Completely dissociate in water, e.g. HCl, HBr, HI, HNO3, H2SO4, HClO4.

Strong Bases

Dissociate completely in water, typically Group 1 and Group 2 hydroxides.

Weak Acids

Partially dissociate in solution; establish an equilibrium.

Ka

Equilibrium constant for weak acids: Ka = [H+][A-]/[HA].

Kb

Equilibrium constant for weak bases: Kb = [HB+][OH-]/[B].

Percent Ionization

% Ionization = [H+]eq / [HA]initial × 100.

Binary Acids

Acids with the formula H-X; acidity increases down a group.

Oxyacids

Acids with the formula H-O-Z; strength depends on electronegativity of Z.

Carboxylic Acids

Acids with a -COOH group where resonance stabilizes the anion.

Neutralization Reaction

Reaction between acid and base that results in water and a salt.

Hydrolysis of Salts

Reactions of ions in solution that can produce acidic or basic solutions.

Buffer Solution

A solution that resists pH changes upon the addition of strong acids/bases.

Henderson-Hasselbalch Equation

pH = pKa + log([A-]/[HA]).

Buffer Capacity

The amount of acid or base a buffer can absorb without significant pH change.

Titration

Controlled neutralization used to determine concentration of an unknown.

Equivalence Point

Point in a titration when moles of acid = moles of base.

Indicators

Weak acids whose color changes when they lose their proton.

Polyprotic Acids

Acids that can donate more than one proton, each with a different Ka.

Spectator Ions

Ions that do not affect pH and do not participate in the reaction.

Ionization and Strength Relationship

Higher Ka means lower pKa and stronger acid.

Le Chatelier’s Principle

If a system at equilibrium is disturbed, it shifts to counteract the disturbance.

Endothermic Reaction

A reaction that absorbs heat, such as the autoionization of water.

Amphoteric Substance

A substance that can act as both an acid and a base.

Hydrolysis

The reaction of a salt with water to produce acidic or basic solutions.

Acidic Salt

Salt derived from a weak base and strong acid; pH < 7 at equivalence.

Basic Salt

Salt derived from a weak acid and strong base; pH > 7 at equivalence.

Common Ion Effect

The shift in equilibrium caused by the addition of a compound that shares an ion.

Molecular Structure Impact

The strength of an acid can depend on the molecular structure and bonding.

Ice Table

A table used to determine the concentrations of species at equilibrium.

Neutral pH

pH of 7 at 25°C; may change with temperature.

Temperature's Effect on Kw

Kw increases with temperature, affecting pH neutrality.

Titration Curve

Graph showing pH change versus the volume of titrant added.

Half-Equivalence Point

Point in titration when concentrations of weak acid and its conjugate base are equal.

Stoichiometry in Titrations

Using stoichiometric relationships to find remaining species in neutralization.

Conjugate Base Stability

Stability of the conjugate base contributes to acid strength.