AP Chemistry Unit 8 Notes

Unit Overview

Unit Title: Acids & Bases

Content Outline:

• 8.1 Introduction to Acids and Bases

• 8.2 pH and pOH of Strong Acids and Bases

• 8.3 Weak Acid and Base Equilibria

• 8.4 Acid-Base Reactions and Buffers

• 8.5 Acid-Base Titrations

• 8.6 Molecular Structure of Acids and Bases

• 8.7 pH and pKa

• 8.8 Properties of Buffers

• 8.9 Henderson-Hasselbalch Equation

• 8.10 Buffer Capacity

• 8.11 pH and Solubility

8.1 Introduction to Acids and Bases

Arrhenius Definition:

• Acids release H+ ions (protons) when dissolved in water, increasing the concentration of hydrogen ions (e.g., HCl → H+ + Cl-).

• Bases release OH- ions in water, increasing the hydroxide ion concentrations (e.g., NaOH → Na+ + OH-).

8.2 pH and pOH of Strong Acids and Bases

Relationships:

• pH is defined as the negative logarithm of the hydrogen ion concentration: pH = -log[H+].

• pOH is determined in a similar manner: pOH = -log[OH-].

• The ion product of water at 25°C is: Kw = [H3O+][OH-] = 1.0 x 10^-14, illustrating the relationship between water's hydrogen and hydroxide ions.

• The relationship between pH and pOH is expressed as: pH + pOH = 14.

Strong Acids:

• In calculations for strong acids, complete ionization is assumed, meaning all acid molecules dissociate in solution.

  • Example: For a 0.20 M HCl solution, it can be calculated that

    • [H+] = 0.20 M

    • pH = 0.70.

Strong Bases:

• Similarly, assumptions of 100% ionization are made for strong bases.

  • Example: For a 0.015 M NaOH solution, it can be calculated that

    • [OH-] = 0.015 M

    • pOH = 1.82, pH = 12.18.

8.3 Weak Acid and Base Equilibria

Equilibrium Expressions:

• Weak acids do not ionize completely; an equilibrium is established as follows:

  • HA ⇌ H+ + A-

• The equilibrium constant for weak acids is denoted as:

  • Ka = [H+][A-]/[HA], reflecting the degree of ionization.

• Henderson-Hasselbalch Equation is utilized to relate pH with pKa:

  • pH = pKa + log([A-]/[HA]).

8.4 Acid-Base Reactions and Buffers

Buffers consist of a weak acid and its conjugate base and act to resist changes in pH by neutralizing small amounts of added acid or base.

• Example: When acetic acid (HC2H3O2) is combined with sodium acetate (NaC2H3O2), it forms a buffer solution that can maintain a relatively stable pH when acids or bases are introduced.

8.5 Acid-Base Titrations

Titration Curves:

• The pH at the equivalence point varies based on the nature of the acid and base involved:

  • Strong Acid/Strong Base: At equivalence, pH = 7.

  • Weak Acid/Strong Base: The equivalence point occurs at pH > 7.

Calculations:

• Stoichiometry is employed to determine the concentrations of reactants and products throughout the titration process.

8.6 Molecular Structure of Acids and Bases

Influence of Structure:

• Electronegativity of atoms within an acid influences its acidity; atoms that have a greater electronegativity tend to pull electron density away from the bond, stabilizing the release of protons.

• Resonance stability in conjugate bases enhances acid strength; for instance,

  • ClO3- is a stronger conjugate base than ClO2- due to resonance stabilizing the negative charge.

8.7 pH and pKa

Concept Overview:

• The lower the pKa value, the stronger the acid; therefore, stronger acids donate protons more readily.

• The relationship between pH and pKa is crucial: when pH < pKa, the acid form predominates; conversely, when pH > pKa, the deprotonated form predominates.

8.8 Properties of Buffers

• Buffers are essential in maintaining a stable pH, enabling biochemical processes to function optimally.

• Capacity:

  • The buffer capacity can be enhanced by increasing the concentrations of the weak acid and its conjugate base within the solution, allowing it to withstand larger volumes of added acids or bases without significant changes in pH.

8.9 Henderson-Hasselbalch Equation

Application:

• This equation assists in the calculation of buffer pH based on the ratio of acid to its conjugate base:

  • pH = pKa + log([base]/[acid]).

8.10 Buffer Capacity

• Buffer capacity is defined as the amount of acid or base a buffer can neutralize before a significant change in pH occurs; it is maximized when the concentrations of the acid and base components are equimolar.

• Testing various ratios of acid and base pairs can validate buffer effectiveness in maintaining pH stability.

8.11 pH and Solubility

• The solubility of salts is affected by pH shifts that influence ionic equilibrium in solution. For example:

  • The introduction of H+ ions can decrease the concentration of basic anions, thus increasing solubility.

• Complex mixtures that involve weak acids and bases can illustrate key principles relevant to solubility predictions, demonstrating how equilibrium shifts directly correlate with acidic or basic environments.