Acids, Bases, and Buffers: General Chemistry Study Notes

Acids and Bases

Introduction to Acids and Bases

Acids and bases are fundamental chemical species that play a crucial role in many chemical reactions and biological processes. Acids are typically characterized by their sour taste, while bases are known for their bitter or salty taste.

• Acids: Substances that produce H3O+ (hydronium ion) in aqueous solution.

• Bases: Substances that produce OH- (hydroxide ion) in aqueous solution.

Arrhenius Definition

The Arrhenius definition is one of the earliest and simplest ways to classify acids and bases.

• Acid: Produces H+ (or H3O+) ions in water.

• Base: Produces OH- ions in water.

Brønsted-Lowry Definition

The Brønsted-Lowry definition expands the concept of acids and bases beyond aqueous solutions.

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

• Conjugate acid-base pair: When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

Examples of Acid-Base Reactions

• CH3COOH(aq) + H2O(l) → CH3COO-(aq) + H3O+(aq)

• NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)

Strong vs. Weak Acids and Bases

• Strong acids/bases: Completely ionize in aqueous solution.

• Weak acids/bases: Partially ionize in aqueous solution.

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water.

• Strong electrolytes: Completely dissociate into ions (e.g., NaCl).

• Weak electrolytes: Partially dissociate (e.g., CH3COOH).

• Nonelectrolytes: Do not dissociate (e.g., distilled water).

Acid and Base Strength

• A strong acid contains a weak conjugate base.

• A strong base contains a weak conjugate acid.

Types of Acids

• Monoprotic acids: Donate one proton (e.g., HCl).

• Diprotic acids: Donate two protons (e.g., H2SO4).

• Triprotic acids: Donate three protons (e.g., H3PO4).

• Amphiprotic: Can act as either an acid or a base (e.g., H2O).

Oxyacids and Organic Acids

• Oxyacids: Acidic hydrogen is attached to an oxygen atom (e.g., H2SO4, H3PO4).

• Organic acids: Contain a carboxyl group (-COOH), usually weak (e.g., CH3COOH).

Naming Acids

• Binary acids: Hydro + base name of nonmetal + -ic acid (e.g., Hydrochloric acid).

• Ternary acids: Based on polyatomic ions: -ate ion → -ic acid, -ite ion → -ous acid (e.g., Nitric acid from nitrate, Nitrous acid from nitrite).

Acid Ionization and pH

Acid Ionization Constant (Ka)

The acid ionization constant, Ka, measures the strength of an acid in solution.

• $Ka = \frac{[A^-][H_3O^+]}{[HA]}$

• For strong acids, Ka is not used because they ionize completely.

• pKa is related to Ka: $pKa = -\log Ka$

• Stronger acids have higher Ka and lower pKa.

Ionization of Water and pH Scale

Water can ionize to form hydronium and hydroxide ions.

• $H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$

• $K_w = [H_3O^+][OH^-] = 1 \times 10^{-14}$ at 25°C

• Neutral solution: $[H^+] = [OH^-]$

• Acidic solution: $[H^+] > [OH^-]$

• Basic solution: $[H^+] < [OH^-]$

pH and pOH Calculations

• $pH = -\log [H_3O^+]$

• $pOH = -\log [OH^-]$

• $pH + pOH = 14$

pH Indicators in Nature

Some plants and flowers change color depending on the pH of the soil.

• Bigleaf hydrangea flowers are pink in acidic soil and blue in basic (alkaline) soil.

Acid-Base Reactions and Neutralization

Neutralization Reactions

Neutralization occurs when an acid reacts with a base to produce a salt and water.

• General equation: Acid + Base → Salt + Water

• Example: $H_2SO_4(aq) + 2NaOH(aq) \rightarrow Na_2SO_4(aq) + 2H_2O(l)$

Titration

Titration is a laboratory technique used to determine the concentration of an unknown acid or base by reacting it with a solution of known concentration.

• Equivalence point: The point at which equal amounts of acid and base have reacted (pH = 7 for strong acid-strong base titrations).

• Calculation: $M_1V_1 = M_2V_2$ (for monoprotic acids and bases)

Buffers

Buffer Solutions

Buffers are solutions that resist changes in pH when small amounts of acid or base are added.

• Composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Example: Acetic acid (CH3COOH) and sodium acetate (CH3COO-Na+).

Buffer Composition and Types

• Carbonate buffer: H2CO3 / HCO3-

• Phosphate buffer: H2PO4- / HPO42-

• Protein buffer: Proteins can act as buffers in biological systems.

Buffer Action

Buffers maintain pH by neutralizing added acids or bases.

• When acid is added: $HCO_3^- + H_3O^+ \rightarrow H_2CO_3 + H_2O$

• When base is added: $H_2CO_3 + OH^- \rightarrow HCO_3^- + H_2O$

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentration of the acid and its conjugate base.

• $pH = pKa + \log \frac{[\text{Conjugate Base}]}{[\text{Weak Acid}]}$

• Used to calculate the pH of buffer solutions.

Summary Table: Acid, Base, and Buffer Properties

Type	Definition	Example
Acid	Donates H+	HCl, CH3COOH
Base	Accepts H+	NaOH, NH3
Buffer	Resists pH change	CH3COOH/CH3COO-, H2CO3/HCO3-

Additional info:

• Blood is a natural buffer, maintaining pH between 7.35 and 7.45.

• pH indicators are used in laboratory and nature to visually determine pH.