Acids, Bases, and Aqueous Equilibria: Study Notes

Acids, Bases, and Aqueous Equilibria

Stomach Acid and Heartburn

Stomach acid, primarily hydrochloric acid (HCl), is produced by cells lining the stomach. Sometimes, this acid can leak into the esophagus, causing discomfort known as heartburn or, in chronic cases, gastroesophageal reflux disease (GERD).

• GERD is caused by the failure of the muscle separating the stomach and esophagus to close tightly, allowing acid to escape.

• Diagnosis often involves attaching a pH sensor to the esophagus to monitor acidity over time.

pH and Related Formulas

The pH scale measures the acidity or basicity of a solution based on the concentration of hydrogen ions.

• pH is defined as: $pH = -\log [H^+]$

• Alternatively, using hydronium ion: $pH = -\log [H_3O^+]$

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

• The ion product of water: $K_w = 1.0 \times 10^{-14} = [H^+][OH^-]$

• To find [H+] from pH: $[H^+] = 10^{-pH}$

• Relationship between pH and pOH: $pH + pOH = 14$

Properties and Structures of Acids

Acids are substances with distinct chemical and physical properties.

• Properties:

  • Sour taste

  • Ability to dissolve many metals

  • Neutralize bases

  • Turn blue litmus paper red

• Structures:

  • Binary acids: Acidic hydrogen attached to a nonmetal atom (e.g., HCl)

  • Oxyacids: Acidic hydrogen attached to an oxygen atom (e.g., HNO3)

  • Carboxylic acids: Contain a –COOH group (e.g., acetic acid, CH3COOH)

Properties of Bases

Bases are substances that exhibit the following properties:

• Bitter taste

• Slippery feel

• Turn red litmus paper blue

• Neutralize acids

Definitions of Acids and Bases

• Arrhenius Definition:

  • Acid: Produces H+ ions in solution

  • Base: Produces OH- ions in solution

• Brønsted-Lowry Definition:

  • Acid: Proton (H+) donor

  • Base: Proton (H+) acceptor

Hydronium Ion and Proton Transfer

• H+ ions are highly reactive and do not exist freely in water; they form hydronium ions (H3O+) by associating with water molecules.

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

• When a base accepts a proton, it becomes a conjugate acid.

Acid and Base Strength

The strength of an acid or base depends on its degree of ionization in water.

• Strong acids: Ionize completely in water; strong electrolytes.

• Strong bases: Form OH- ions completely; strong electrolytes.

• Weak acids: Only partially ionize; weak electrolytes.

• Weak bases: Only partially form OH- ions; weak electrolytes.

Examples of Strong Acids

• Hydrochloric acid (HCl)

• Nitric acid (HNO3)

• Hydrobromic acid (HBr)

• Perchloric acid (HClO4)

• Hydroiodic acid (HI)

• Sulfuric acid (H2SO4) (diprotic)

Examples of Weak Acids

• Acetic acid (CH3COOH)

• Formic acid (HCOOH)

• Phosphoric acid (H3PO4)

• Carbonic acid (H2CO3)

• Most organic acids

• Additional info: Only a small fraction of weak acid molecules donate H+ to water.

Acid Ionization Constant (Ka)

The acid ionization constant quantifies the strength of an acid in water.

• Larger Ka = stronger acid

• Expression: $K_a = \frac{[H_3O^+][A^-]}{[HA]} = \frac{[H^+][A^-]}{[HA]}$

Autoionization of Water

Water can act as both an acid and a base (amphoteric), leading to a small degree of self-ionization.

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

• At 25°C: $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7} \text{ M}$

• Ion product: $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

pH, pOH, and Solution Classification

• Neutral solution: $[H_3O^+] = [OH^-]$

• Acidic solution: $[H_3O^+] > [OH^-]$

• Basic solution: $[OH^-] > [H_3O^+]$

• pH scale:

  • pH < 7: acidic

  • pH = 7: neutral

  • pH > 7: basic

• pOH is another way to express solution basicity: $pOH = -\log [OH^-]$

• Relationship: $pH + pOH = 14.00$

pKa and pKb

• pKa and pKb are logarithmic measures of acid and base strength.

• The stronger the acid, the smaller the pKa.

• The stronger the base, the smaller the pKb.

[H3O+] and pH of a Strong Acid

• In a strong acid solution, [H3O+] comes almost entirely from the acid, not from water autoionization.

• Strong acids dissociate completely (single arrow, not equilibrium).

• Stoichiometry is straightforward: the concentration of the acid equals the concentration of H3O+.

Acid-Base Properties of Ions and Salts

Salts, when dissolved in water, can affect the pH depending on the nature of their constituent ions.

• Salts of strong acids and strong bases: Do not affect pH (neutral solutions).

• Salts with cations of weak bases or anions of weak acids: Can hydrolyze and affect pH.

• Anions: Tend to form basic or neutral solutions.

• Cations: Tend to form acidic or neutral solutions.

Anions as Weak Bases

• Every anion is the conjugate base of an acid.

• Anions from strong acids are pH-neutral (do not react with water).

• Anions from weak acids are basic (can accept protons from water).

• The stronger the acid, the weaker its conjugate base.

Relationship Between Ka, Kb, and Kw

• For a conjugate acid-base pair: $K_a \times K_b = K_w$

• The stronger the base, the weaker its conjugate acid, and vice versa.

Summary Table: Acid and Base Strengths

Type	Degree of Ionization	Electrolyte Strength	Example
Strong Acid	Nearly 100%	Strong	HCl, HNO3
Weak Acid	Small fraction	Weak	CH3COOH
Strong Base	Nearly 100%	Strong	NaOH, KOH
Weak Base	Small fraction	Weak	NH3

Example: Calculating pH

• Given [H+] = 1.0 × 10-3 M, calculate pH:

$pH = -\log(1.0 \times 10^{-3}) = 3$

Example: Relationship of Ka and Kb

• If the Ka of acetic acid is 1.8 × 10-5, what is the Kb of its conjugate base?

$K_b = \frac{K_w}{K_a} = \frac{1.0 \times 10^{-14}}{1.8 \times 10^{-5}} = 5.6 \times 10^{-10}$

Additional info: For weak bases, always include water as a reactant in the equilibrium equation.