Acids and Bases

10.1 Acids and Bases: Definitions

Acids and bases are fundamental chemical species that participate in a wide variety of chemical reactions. Their definitions have evolved to encompass different chemical contexts.

• Arrhenius Definition:

  • Acid: A substance that produces hydronium ions (H3O+) when dissolved in water.

  • Base: A substance that produces hydroxide ions (OH-) when dissolved in water.

  • Example: HCl (acid) in water forms H3O+; NaOH (base) in water forms OH-.

• Brønsted-Lowry Definition:

  • Acid: Proton (H+) donor.

  • Base: Proton (H+) acceptor.

  • This definition is broader and not limited to aqueous solutions.

  • Example: NH3 (ammonia) acts as a base by accepting a proton from water to form NH4+.

• Monoprotic, Diprotic, and Triprotic Acids:

  • Monoprotic: Can donate one proton (e.g., HCl).

  • Diprotic: Can donate two protons (e.g., H2SO4).

  • Triprotic: Can donate three protons (e.g., H3PO4).

• Acidic Hydrogens: Acidic H atoms are typically bonded to electronegative atoms (e.g., O, N, or halogens).

• 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.

  • Example: HCN (acid) and CN- (conjugate base); HClO4 (acid) and ClO4- (conjugate base).

10.2 Acid and Base Strength

The strength of acids and bases is determined by their ability to donate or accept protons.

• Strong Acids: Completely dissociate in water (e.g., HCl, HNO3).

• Weak Acids: Partially dissociate in water (e.g., CH3COOH).

• Strong Bases: Completely dissociate in water (e.g., NaOH, KOH).

• Weak Bases: Partially accept protons (e.g., NH3).

• Polyprotic Acids: Undergo stepwise dissociation; each successive proton is harder to remove.

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

• Equilibrium Direction: Proton-transfer equilibria favor the formation of the weaker acid and base.

Table: Common Acids and Bases (Summary)

10.3 Acid Dissociation Constants

The extent to which a weak acid dissociates in water is quantified by its acid dissociation constant, Ka.

• Acid Dissociation Constant (Ka):

  • For a generic acid HA:

• Interpretation:

  • Strong acids: (dissociation favored)

  • Weak acids: (dissociation not favored)

  • Each successive dissociation in polyprotic acids has a lower Ka.

10.4 Water as Both an Acid and a Base

Water is amphoteric, meaning it can act as either an acid or a base depending on the reaction partner.

• Self-Ionization of Water:

  • Ion-product constant: at 25°C

• Acidic, Neutral, and Basic Solutions:

  • Acidic:

  • Neutral:

  • Basic:

10.5 Measuring Acidity in an Aqueous Solution: The pH Scale

The pH scale is a logarithmic measure of the hydronium ion concentration in solution.

• pH Definition:

• pOH Definition:

• Relationship: at 25°C

• Interpretation:

  • pH < 7: Acidic

  • pH = 7: Neutral

  • pH > 7: Basic

• Indicators: Dyes that change color depending on pH (e.g., litmus, phenolphthalein).

10.6 Working with pH

Conversions between pH and hydronium ion concentration are common in acid-base chemistry.

• From pH to [H3O+]:

• From [H3O+] to pH:

• Significant Figures: The number of decimal places in pH equals the number of significant figures in [H3O+].

10.7 Acid and Base Equivalents

Equivalents and normality are used to express the reactive capacity of acids and bases.

• Equivalent (Eq):

  • For acids: Amount containing 1 mole of H+ ions.

  • For bases: Amount containing 1 mole of OH- ions.

• Normality (N): Number of equivalents per liter of solution.

• Relationship: (where n = number of H+ or OH- ions per formula unit)

• Example: For H2SO4 (diprotic acid), 1 mole yields 2 equivalents.

10.8 Some Common Acid-Base Reactions

Acids and bases participate in several important types of reactions, including neutralization and reactions with carbonates and amines.

• Neutralization (Acid + Base):

  • General:

  • Example:

• Acid + Carbonate/Bicarbonate:

  • Used in antacids and geological tests for carbonates.

• Acid + Ammonia (or Amines):

  • Produces water-soluble ammonium salts.

10.9 Acidity and Basicity of Salt Solutions

The pH of a salt solution depends on the strengths of the acid and base from which the salt is derived.

• Salt from Strong Acid + Strong Base: Neutral solution (e.g., NaCl).

• Salt from Strong Acid + Weak Base: Acidic solution (e.g., NH4Cl).

• Salt from Weak Acid + Strong Base: Basic solution (e.g., NaCN, NaHCO3).

• Both ions react with water: The ion that reacts more strongly determines the pH.

Table: Acidity and Basicity of Salt Solutions (Summary)

10.10 Buffer Solutions

Buffers are solutions that resist changes in pH when small amounts of acid or base are added. They are essential in biological and chemical systems.

• Composition: Usually a weak acid and its conjugate base (or weak base and its conjugate acid).

• Action:

  • Added acid is neutralized by the base component.

  • Added base is neutralized by the acid component.

• Henderson-Hasselbalch Equation:

  • Where [A-] is the conjugate base and [HA] is the weak acid.

• Effective Buffering:

  • pKa of the acid should be close to the desired pH.

  • The ratio [A-]/[HA] should be close to 1.

  • The buffer components should be present in much greater amounts than the acid or base to be neutralized.

10.11 Titration

Titration is a laboratory technique used to determine the concentration of an acid or base in solution.

• Process:

  • A solution of known concentration (titrant) is added to a measured volume of solution of unknown concentration until the reaction is complete (end point, indicated by a color change).

• Calculations:

  • For monoprotic acids and bases:

  • For polyprotic acids or bases, or when stoichiometry is not 1:1, use equivalents and normality:

• Example: If 44.5 mL of 0.100 M NaOH is required to neutralize 5.00 mL of vinegar, the molarity of acetic acid can be calculated using the above relationships.

Additional info: Where the original text was incomplete or missing, standard chemical knowledge and textbook conventions were used to fill in definitions, equations, and examples. Tables were reconstructed based on typical content for GOB Chemistry.