2.2 Acid-base chemistry and pH

Acid-base chemistry is all about proton exchange. It's crucial for understanding how molecules interact in solution and how pH affects biological systems. This topic covers different theories of acids and bases, from Arrhenius to Brønsted-Lowry.

pH and pOH measure acidity and basicity in solutions. The pH scale, dissociation constants, and acid-base equilibria help us predict and control chemical reactions in water. These concepts are vital for understanding buffer systems in living organisms.

Acid-Base Theories

Arrhenius Theory

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• Defines acids as substances that dissociate in water to produce hydrogen ions (H+)
• Bases dissociate in water to produce hydroxide ions (OH-)
• Limited in scope as it only applies to aqueous solutions and does not account for reactions in other solvents or gases

Brønsted-Lowry Theory

• Defines acids as proton (H+) donors
• Bases are proton acceptors
• Extends the concept of acids and bases beyond aqueous solutions
• Accounts for acid-base reactions in various solvents and gas-phase reactions
• Ammonia (NH3) acts as a base by accepting a proton to form ammonium ion (NH4+)

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
• Conjugate acid-base pairs differ by a single proton (H+)
• Examples:
  • Acetic acid (CH3COOH) and acetate ion (CH3COO-)
  • Ammonia (NH3) and ammonium ion (NH4+)

Acid and Base Strength

• Strong acids and bases completely dissociate in water
• Weak acids and bases only partially dissociate in water
• Strength depends on the extent of dissociation
• Examples of strong acids: Hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3)
• Examples of weak acids: Acetic acid (CH3COOH), carbonic acid (H2CO3)
• Examples of strong bases: Sodium hydroxide (NaOH), potassium hydroxide (KOH)
• Examples of weak bases: Ammonia (NH3), methylamine (CH3NH2)

pH and pOH

pH Scale

• Measures the acidity or basicity of a solution
• Ranges from 0 to 14
• Neutral solutions have a pH of 7
• Acidic solutions have a pH below 7
• Basic solutions have a pH above 7
• Each unit change in pH represents a tenfold change in H+ concentration

pOH

• Measures the hydroxide ion concentration in a solution
• Related to pH by the equation: pH + pOH = 14
• Neutral solutions have a pOH of 7
• Acidic solutions have a pOH above 7
• Basic solutions have a pOH below 7

Dissociation Constant (Ka)

• Quantifies the strength of an acid
• Represents the equilibrium constant for the dissociation of a weak acid
• Higher Ka values indicate stronger acids
• Ka = [H+][A-] / [HA], where HA is the weak acid, A- is its conjugate base
• Example: Acetic acid (CH3COOH) has a Ka of 1.8 × 10^-5

pKa

• Negative logarithm of the acid dissociation constant (Ka)
• pKa = -log(Ka)
• Higher pKa values indicate weaker acids
• Useful for comparing acid strengths
• Example: Acetic acid (CH3COOH) has a pKa of 4.74

Acid-Base Equilibria

Henderson-Hasselbalch Equation

• Relates pH to the pKa and the ratio of the concentrations of a weak acid and its conjugate base
• pH = pKa + log([A-] / [HA])
• Useful for calculating the pH of buffer solutions
• Helps predict the pH changes during titrations
• Example: Calculate the pH of a buffer solution containing 0.1 M acetic acid (CH3COOH) and 0.2 M sodium acetate (CH3COONa), given the pKa of acetic acid is 4.74

Titration

• Technique used to determine the concentration of an unknown acid or base
• Involves the gradual addition of a standard solution (titrant) to the unknown solution (analyte)
• Endpoint is reached when the reaction between the titrant and analyte is complete
• Indicated by a color change in the presence of an indicator or a sharp change in pH
• Types of titrations:
  • Acid-base titration
  • Redox titration
  • Complexometric titration
• Titration curves show the pH changes during the titration process
• Equivalence point is the point at which the moles of titrant added equal the moles of analyte present