Acids and Bases Notes

Acids and Bases Notes

Conjugate Acid-Base Pairs

  • Definition: Conjugate acid-base pairs consist of species that differ by a single proton.

    • Example Reaction:

    • HX + H2O ⇌ X− + H3O+

  • Identification:

    • Acid: Species donating the proton (e.g., HX)

    • Base: Species accepting the proton (e.g., X−)

Relative Strengths of Acids and Bases

  • Important Points:

    • The conjugate base of a strong acid is a very weak base (negligible basicity).

    • The conjugate base of a weak acid is a weak base.

    • The conjugate base of a very weak acid is a strong base.

  • Equilibrium Favors:

    • The reaction tends to favor the formation of weaker acids and bases.

    • Example Reaction:

    • HSO4− + CO32− ⇌ SO42− + HCO3−

Autoionization of Water

  • Water can act both as an acid and a base, resulting in its self-ionization:

    • Reaction: H2O ⇌ H3O+ + OH−

  • Important Characteristics:

    • Rapid equilibrium where water does not stay ionized for long.

    • Kw = [H3O+][OH−] = 1.0 x 10−14 at 25°C

The pH Scale

  • pH Calculation:

    • Formula: pH = -log[H+]

  • Solution Classification:

    • Acidic: [H+] > 1.0 x 10−7 M; pH < 7

    • Neutral: [H+] = 1.0 x 10−7 M; pH = 7

    • Basic: [H+] < 1.0 x 10−7 M; pH > 7

The pOH Scale

  • pOH Calculation:

    • Formula: pOH = -log[OH−]

  • Relationship: pH + pOH = 14.00

  • Example: Calculate [H+] given a pOH of 10.24.

Measuring pH: Indicators

  • Indicators change color at different pH levels (Color Change Ranges):

    • Methyl violet: Yellow (pH 4) to violet (pH 6)

    • Phenolphthalein: Colorless (pH < 8) to pink (pH > 8)

Strong Acids and Bases

  • Strong Acids: Completely dissociate in solution (e.g., HCl, H2SO4).

    • No equilibrium arrow in equations owing to 100% dissociation.

  • Strong Bases: Hydroxide compounds that also fully dissociate (e.g., NaOH, Ca(OH)2).

Weak Acids

  • A weak acid establishes an equilibrium, partially dissociating in water:

    • Reaction: HA(aq) + H2O ⇌ H3O+(aq) + A−(aq)

  • Ka: Acid-dissociation constant, higher value indicates stronger acid.

  • Example: Given [H+] and equilibrium parameters, calculate Ka.

Polyprotic Acids

  • Acids with multiple ionizable protons (e.g., H2SO4).

    • Ka values decrease for each successive dissociation.

    • If Ka1 > Ka2 by a factor of 1000, treat the acid as monoprotic for pH calculations.

Percent Ionization

  • Formula: % Ionization = [H+]eq / [HA]initial x 100%

  • Use to measure strength of weak acids or weak bases.

Buffers

  • Solutions that resist changes in pH upon the addition of small amounts of acid or base.

  • Made using a weak acid/base and its conjugate.

  • Henderson-Hasselbalch Equation: pH = pKa + log([A−]/[HA])

Titrations

  • Titration Steps:

    1. Initial pH determined by the acid/base.

    2. pH changes calculated using stoichiometry until equivalence point (where moles of acid = moles of base).

    3. At equivalence point for strong acid-strong base, pH = 7.

    4. For weak acid-strong base, pH > 7 due to the formation of a weak base.

Relationships in Acid-Base Solutions

  • Ka and Kb Relationship:

    • Ka × Kb = Kw

    • As acid strength increases, conjugate base strength decreases.

Common Ion Effect

  • The presence of a common ion reduces the dissociation of weak acids or bases in a solution.

    • Example: Adding a salt to a weak acid shifts equilibrium left.

Anions and Cations in Solution

  • Cations from weak acids can lead to acidic solutions; anions from weak bases can cause basic solutions.

  • If both react with water, the resultant solution's pH depends on their relative strengths.

Practice Problems

  • Calculate the pH of various solutions, determine Ka and Kb values, and analyze titrations with the Henderson-Hasselbalch equation.