Aqueous Equilibria: Acids, Bases, and Buffers

Acids and Bases: General Properties and Definitions

Introduction to Acids and Bases

Acids and bases are fundamental classes of compounds in chemistry, characterized by their distinct chemical behaviors and roles in aqueous equilibria. Understanding their properties, definitions, and reactions is essential for mastering general chemistry.

Stomach Acid and Heartburn

• Hydrochloric acid (HCl) is produced by cells lining the stomach.

• Functions of stomach acid include:

  • Killing unwanted bacteria

  • Breaking down food

  • Activating digestive enzymes

• If stomach acid backs up into the esophagus, it causes irritation known as heartburn or acid reflux. Chronic cases are referred to as GERD (Gastroesophageal Reflux Disease).

General Properties of Acids

• Acids are defined by their chemical reactivity, particularly their ability to donate protons (H+).

• Common properties:

  • Sour taste

  • Ability to dissolve many metals

  • Neutralize bases

  • Turn blue litmus paper red

• Examples: Acetic acid (in vinegar), hydrochloric acid (in stomach acid).

General Properties of Bases

• Bases are defined by their ability to accept protons or produce hydroxide ions (OH−) in solution.

• Common properties:

  • Bitter taste

  • Slippery feel

  • Neutralize acids

  • Turn red litmus paper blue

• Examples: Sodium hydroxide (NaOH), ammonia (NH3).

Acid and Base Definitions

Arrhenius Definition

• Arrhenius acid: Produces H+ ions in aqueous solution.

• Arrhenius base: Produces OH− ions in aqueous solution.

• Limitation: Does not account for all acid-base reactions, especially those not involving water or hydroxide ions.

Brønsted–Lowry Definition

• Brønsted–Lowry acid: Proton (H+) donor.

• Brønsted–Lowry base: Proton (H+) acceptor.

• This definition is broader and includes reactions outside aqueous solutions.

• All Arrhenius acids and bases are also Brønsted–Lowry acids and bases.

Conjugate Acid–Base Pairs

• When an acid donates a proton, it forms its conjugate base.

• When a base accepts a proton, it forms its conjugate acid.

• Every acid–base reaction involves two conjugate pairs.

Acid and Base Strength

Strong and Weak Acids/Bases

• Strong acids and bases completely dissociate in water.

• Weak acids and bases only partially dissociate, establishing an equilibrium.

• Examples of strong acids: HCl, HNO3, H2SO4.

• Examples of weak acids: HF, CH3COOH.

Quantifying Acid and Base Strength

• The acid dissociation constant ($K_a$) measures the strength of an acid:

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

• The base dissociation constant ($K_b$) measures the strength of a base:

$K_b = \frac{[BH^+][OH^-]}{[B]}$

• The larger the $K_a$ or $K_b$, the stronger the acid or base.

• p$K_a$ and p$K_b$ are logarithmic measures:

$pK_a = -\log K_a$ $pK_b = -\log K_b$

• The smaller the p$K_a$ or p$K_b$, the stronger the acid or base.

Autoionization of Water and the pH Scale

• Water is amphoteric and can act as both an acid and a base.

• Autoionization: $2H_2O \rightleftharpoons H_3O^+ + OH^-$

• The ion product constant for water: $K_w = [H^+][OH^-] = 1.00 \times 10^{-14}$ at 25°C.

• pH is defined as $pH = -\log[H^+]$

• pOH is defined as $pOH = -\log[OH^-]$

• Relationship: $pH + pOH = 14$

pH of Common Substances

• pH values of everyday substances range from highly acidic (gastric juice, pH 1–3) to highly basic (household ammonia, pH 11–12).

Buffer Solutions

Definition and Function

A buffer solution resists changes in pH when small amounts of acid or base are added. Buffers are typically composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Buffer capacity depends on the ratio and concentration of the acid/base pair.

• Blood plasma is an example of a natural buffer system (H2CO3/HCO3−).

Henderson–Hasselbalch Equation

• The pH of a buffer can be calculated using the Henderson–Hasselbalch equation:

$pH = pK_a + \log \left( \frac{[A^-]}{[HA]} \right)$

• This equation is valid when the concentrations of acid and conjugate base are much greater than the $K_a$ value.

Titration and Titration Curves

Acid–Base Titration

In a titration, a solution of known concentration (titrant) is added to a solution of unknown concentration until the reaction reaches the equivalence point, where stoichiometric amounts of acid and base have reacted.

• An indicator is used to signal the endpoint by a color change.

Titration Curves

• Titration curves plot pH versus volume of titrant added.

• The inflection point corresponds to the equivalence point.

• The pH at equivalence depends on the nature of the acid and base:

  • Strong acid + strong base: pH = 7 at equivalence

  • Weak acid + strong base: pH > 7 at equivalence

  • Weak base + strong acid: pH < 7 at equivalence

Summary Table: Acid–Base Properties of Salts

Anion of a	Cation of a	Aqueous Solutions Are	Example
Strong acid	Strong base	Neutral	NaCl
Strong acid	Weak base	Acidic	NH4Cl
Weak acid	Strong base	Basic	NaF
Weak acid	Weak base	Neutral, acidic, or basic (depends on $K_a$ and $K_b$)	NH4F

Key Equations

• $K_w = [H^+][OH^-] = 1.00 \times 10^{-14}$

• $pH = -\log[H^+]$

• $pOH = -\log[OH^-]$

• $pH + pOH = 14$

• $pK_a = -\log K_a$

• $pK_b = -\log K_b$

• $pK_w = pK_a + pK_b$

• Henderson–Hasselbalch: $pH = pK_a + \log \left( \frac{[A^-]}{[HA]} \right)$