Comprehensive Overview of Acids and Bases in Honors Chemistry

Introduction to Acids and Bases

Acids and bases can be defined using different definitions depending on their properties and the reaction types considered. In general inorganic chemistry contexts, the Bronsted-Lowry definitions are prominently used:

Acid: A substance that dissociates to release H⁺ ions (proton donor).
Base: A substance that accepts H⁺ ions (proton acceptor).

Important notes include:

When an acid dissociates in water, the hydrogen ion effectively consists only of a proton.
The dissociation of H⁺ leads to the formation of the hydronium ion (H₃O⁺), since H⁺ readily bonds with water.

It is important to note that in solutions containing H⁺ ions, we can use the equality: [H⁺] = [H₃O⁺]

Strong vs. Weak Acids

Definitions

Understanding whether an acid is strong or weak is crucial:

Strong Acids: Completely dissociate into ions in solution. Example:

HNO₃(aq) + H₂O(l) → NO₃^-(aq) + H₃O⁺(aq)

Weak Acids: Partially dissociate, establishing an equilibrium in solution. Example:

HF(aq) + H₂O(l) ↔ F^-(aq) + H₃O⁺(aq)

Common strong acids to memorize include:

HCl - Hydrochloric acid
HBr - Hydrobromic acid
HI - Hydriodic acid
HNO₃ - Nitric acid
H₂SO₄ - Sulfuric acid (first H⁺ only; the second is weak)
HClO₃ - Chloric acid
HClO₄ - Perchloric acid

Strong vs. Weak Bases

Definitions

Identifying strong and weak bases is generally simpler:

Strong Bases: Soluble hydroxide salts, with examples like NaOH and KOH.
Weak Bases: Typically organic compounds called amines, characterized by an R-NH₂ structure where R can be hydrogen or hydrocarbon groups like CH₃.

Examples of weak bases include ammonia (NH₃) and methyl amine (CH₃NH₂).

Naming Acids

Categories of Acids

Acids can be classified into two major categories for naming:

Binary Acids: Composed of hydrogen and a single nonmetal anion.

Nomenclature: Prefix “hydro” + anion name (change “ide” to “ic”) + “acid”.
Examples:

HCl - Hydrochloric acid
H₂S - Hydrosulfuric acid
HCN - Hydrocyanic acid

Oxoacids: Consist of hydrogen bonded to oxygen-containing polyatomic anions.

No “hydro” prefix used; nomenclature is based on the polyatomic ion name.
For oxoanions:

“ate” becomes “ic”
“ite” becomes “ous”

Examples of Naming Oxoacids

H₂SO₃ is sulfurous acid (SO₃^2- is sulfite).
H₂SO₄ is sulfuric acid.

Acid-Base Reactions

In acid-base reactions, the acid donates H⁺ ions to the base. The products can be determined by swapping hydrogen and adjusting charges, considering the following combinations:

Strong Acid + Strong Base
Strong Acid + Weak Base
Weak Acid + Strong Base
Weak Acid + Weak Base

Key considerations during these reactions include:

Strong reactants lead to one-way reactions.
Every strong reactant contributes spectator ions to the overall reaction.

Examples of Reacting Combinations

Strong Acid + Strong Base:

Example: HCl + KOH → KCl + H₂O
Net Ionic Equation: H⁺(aq) + OH^-(aq) → H₂O(l)

Strong Acid + Weak Base:

Example: HNO₃ + NH₃ → NO₃^-(aq) + NH₄⁺(aq)

Weak Acid + Strong Base:

Example: HF + NaOH → NaF + H₂O
Net Ionic Equation: HF(aq) + OH^-(aq) → H₂O(l) + F^-(aq)

Weak Acid + Weak Base:

Example: HC₂H₃O₂ + CH₃NH₂ ⇌ C₂H₃O₂^- + CH₃NH₃⁺

Stoichiometry Involving Strong Acids and Bases

In strong acid or base reactions, stoichiometry is focused on achieving an exact molar ratio, termed the equivalence point, which is often determined through titration.

Key apparatus for titrations include:

Burette: Holds the titrant (base) with a tap for precise dispensing.
Erlenmeyer Flask: Receives the analyte (acid).

Indicators, like phenolphthalein, signal the endpoint of the reaction:

Color change from colorless in acidic conditions to pink in basic conditions.

pH, pOH, and Equilibrium

The pH scale is vital for determining whether a substance is acidic, basic, or neutral in solution:

pH < 7 indicates acidity.
pH = 7 indicates neutrality.
pH > 7 indicates basicity.

Measuring pH

pH and pOH are related, with the following equations governing their relationships:

pH = -log[H₃O⁺]
pOH = -log[OH^-]
[H₃O⁺][OH^-] = 1.0 x 10^-14

Thus, for any aqueous solution:

pH + pOH = 14

Equilibrium of Water

Water (H₂O) behaves as an amphoteric substance, meaning it can act either as an acid or a base:

H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH^-(aq)

The equilibrium constant for this reaction (K_w) is 1.0 x 10^-14.

Weak Acid and Base Equilibria

Weak acids and bases establish equilibria when dissolved in water:

Weak acid: HNO₂(aq) + H₂O(l) ⇌ NO₂^-(aq) + H₃O⁺(aq)

In these reactions:

Identify the acid, base, conjugate acid, and conjugate base.

Equilibrium Applications

Weak acids (K_a) and weak bases (K_b) have equilibrium constants. Their relationship is described as follows:

K_a x K_b = K_w

Where K_w is 1.00 x 10^-14 at 25 ^oC.

The Acid-Base Nature of Salts

Salts can also exhibit acidic or basic properties in solution:

Analyze each ion in the salt for acidic or basic characteristics.
Example of Acidic Salt: NH₄Br: NH₄⁺ loses H⁺ creating NH₃, a weak base.
Example of Basic Salt: NaNO₂: NO₂^- functions as a base.

Titration Curves and pH at Equivalence

Titration curves illustrate how pH changes as an acid-base reaction progresses:

A strong acid with a strong base results in a neutral salt and a pH of 7 at equivalence (Example: HCl + NaOH).
A weak acid titrated with a strong base results in a higher pH due to the basicity of the conjugate base (Example: acetic acid with NaOH).
A strong acid titrated with a weak base yields a low pH at equivalence due to the acidic character of the resulting cation.

Practice Problems

Write balanced equations for titrations and predict pH at equivalence based on the strength of the acid or base involved, such as:

KOH into HNO₃.
NaOH into HOCl.
Calculate pH at equivalence for HCN titrated with NaOH.