Acids and Bases: Properties, Definitions, and Calculations

Acids and Bases

Introduction to Acids and Bases

Acids and bases are fundamental classes of compounds in chemistry, playing crucial roles in chemical reactions, biological systems, and industrial processes. Their properties, definitions, and behaviors are central to understanding chemical equilibrium and solution chemistry.

Properties and Examples of Acids

General Properties of Acids

• Sour taste (e.g., citric acid in lemons)

• Ability to dissolve many metals (e.g., HCl reacts with Zn to produce H2 gas)

• Ability to neutralize bases

• Change blue litmus paper to red

Common Acids and Their Uses

Name	Formula	Occurrence/Uses
Hydrochloric acid	HCl	Stomach acid, metal cleaning, food preparation
Sulfuric acid	H2SO4	Fertilizer, batteries, dye production
Nitric acid	HNO3	Fertilizer, explosives, dyes
Acetic acid	HC2H3O2	Vinegar, food preservation
Citric acid	H3C6H5O7	Citrus fruits, pH adjustment in foods
Phosphoric acid	H3PO4	Fertilizer, beverages, buffering

Structures of Acids

• Binary acids: H attached to a nonmetal (e.g., HCl, HF)

• Oxyacids: H attached to an oxygen atom (e.g., H2SO4, HNO3)

• Carboxylic acids: Contain the COOH group (e.g., acetic acid, citric acid)

Properties and Examples of Bases

General Properties of Bases

• Bitter taste (e.g., alkaloids in plants)

• Slippery feel (e.g., soap)

• Turn red litmus paper blue

• Ability to neutralize acids

Common Bases and Their Uses

Name	Formula	Occurrence/Uses
Sodium hydroxide	NaOH	Soap, plastic, petroleum processing
Potassium hydroxide	KOH	Soap, batteries, cotton processing
Sodium bicarbonate	NaHCO3	Baking soda, antacid
Ammonia	NH3	Detergent, fertilizer, cleaning agent

Definitions of Acids and Bases

Arrhenius Definition

• Acid: Produces H+ ions in aqueous solution

• Base: Produces OH− ions in aqueous solution

Hydronium Ion Formation

In water, H+ ions associate with H2O to form the hydronium ion (H3O+):

$\mathrm{H^+ + H_2O \rightarrow H_3O^+}$

Brønsted–Lowry Definition

• Acid: Proton (H+) donor

• Base: Proton (H+) acceptor (must have a lone pair)

All Arrhenius acids/bases are also Brønsted–Lowry acids/bases, but the Brønsted–Lowry definition is broader and applies to more reactions.

Lewis Definition

• Acid: Electron pair acceptor

• Base: Electron pair donor

This definition is the most general and includes many reactions not covered by the other definitions.

Acid–Base Reactions and Conjugate Pairs

Acid–Base Reactions

In an acid–base reaction, a proton is transferred from the acid to the base. The products are a conjugate base and a conjugate acid.

• Conjugate acid: The species formed when a base gains a proton

• Conjugate base: The species formed when an acid loses a proton

Amphoteric Substances

Amphoteric substances can act as either an acid or a base. Water is the most common example:

$\mathrm{H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-}$

Strength of Acids and Bases

Strong vs. Weak Acids and Bases

• Strong acids/bases: Completely ionize in water (strong electrolytes)

• Weak acids/bases: Partially ionize in water (weak electrolytes)

Examples of Strong and Weak Acids

Strong Acids	Weak Acids
HCl, HNO3, H2SO4, HBr, HI, HClO4	HF, H2CO3, H3PO4, CH3COOH

Acid Ionization Constant (Ka)

The strength of an acid is measured by its acid ionization constant, Ka:

$\mathrm{HA + H_2O \rightleftharpoons A^- + H_3O^+}$

$\mathrm{K_a = \frac{[A^-][H_3O^+]}{[HA]}}$

Larger Ka = stronger acid; smaller Ka = weaker acid.

Autoionization of Water and the pH Scale

Autoionization of Water

Water can ionize to form hydronium and hydroxide ions:

$\mathrm{2H_2O \rightleftharpoons H_3O^+ + OH^-}$

The ion product constant for water is:

$\mathrm{K_w = [H_3O^+][OH^-] = 1.00 \times 10^{-14}}$ at 25°C

pH and pOH

• pH = -\log[H_3O^+]

• pOH = -\log[OH^-]

• pH + pOH = 14.00 (at 25°C)

• [H3O+] = 10^{-pH}

Interpreting pH Values

• pH < 7: Acidic solution

• pH = 7: Neutral solution

• pH > 7: Basic solution

• Each pH unit represents a tenfold change in [H3O+]

pKa and pKb

• pKa = -\log Ka

• pKb = -\log Kb

• Stronger acids have smaller pKa values; stronger bases have smaller pKb values.

Sample Calculations

• Given [H3O+] = 9.2 × 10−9 M, pH = −log(9.2 × 10−9) = 8.04

• Given pH = 8.37, [H3O+] = 10−8.37 = 4.3 × 10−9 M

Summary Table: Acid–Base Concepts

Concept	Definition	Key Equation
Arrhenius Acid	Produces H+ in water	HCl(aq) → H+(aq) + Cl−(aq)
Arrhenius Base	Produces OH− in water	NaOH(aq) → Na+(aq) + OH−(aq)
Brønsted–Lowry Acid	Proton donor	HA + H2O → A− + H3O+
Brønsted–Lowry Base	Proton acceptor	B + H2O → HB+ + OH−
Lewis Acid	Electron pair acceptor	—
Lewis Base	Electron pair donor	—