Acids and Bases Chemistry Notes

Acids and Bases Chemistry Notes

Brønsted-Lowry Acids and Bases

  • Brønsted-Lowry acids: Can transfer a proton ($H^+$) to another substance.

    • Example:
      HCl(aq)+H<em>2O(l)⇌H</em>3O+(aq)+Cl−(aq)HCl(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + Cl^-(aq)

  • Brønsted-Lowry bases: Can accept a proton.

    • Example:
      NH<em>3(aq)+H</em>2O(l)⇌NH4+(aq)+OH−(aq)NH<em>3(aq) + H</em>2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)

  • A Brønsted-Lowry base must have at least one unshared pair of electrons available for bonding with the proton.

  • Note: Water can act as both an acid and a base. The reverse reaction can also be explored, leading to conjugate pairs.

Conjugate Acid-Base Pairs
  • Conjugate acid-base pairs: Species that differ by only one proton ($H^+$).
    • Important: Consider the change in charge when writing them.

Dissociation of Water

  • Water can react as an acid or base:

    • Auto-ionization of water occurs, though only to a minor extent (pure water at 25 °C:
      [H3O+]=1.0×10−7 M[H_3O^+] = 1.0 \times 10^{-7} \text{ M} and
      [OH−]=1.0×10−7 M[OH^-] = 1.0 \times 10^{-7} \text{ M}).

  • Ion-product constant of water:
    K<em>w=[H</em>3O+][OH−]=1.0×10−14.K<em>w = [H</em>3O^+][OH^-] = 1.0 \times 10^{-14}.

Hydronium Ion Concentration and Acidity

  • Acidity classification based on hydronium ion concentration:
    • Acidic: [H<em>3O+]>[OH−][H<em>3O^+] > [OH^-] ; [H3O^+] > 1.0 \times 10^{-7} ext{ M}
    • Neutral: [H3O+]=[OH−]=1.0×10−7extM[H_3O^+] = [OH^-] = 1.0 \times 10^{-7} ext{ M}
    • Basic: [H<em>3O+]<[OH−][H<em>3O^+] < [OH^-] ; [H3O^+] < 1.0 \times 10^{-7} ext{ M}

The pH Scale

  • The logarithmic scale simplifies dealing with small $[H_3O^+]$ values by using the notation $pH$:
    • pH=−log⁡([H3O+]).pH = -\log([H_3O^+]).
  • Examples:
    • Calculate $pH$ from $[H_3O^+]$:
    • If [H3O+]=1.0×10−5extM[H_3O^+] = 1.0 \times 10^{-5} ext{ M}, then pH=−5.pH = -5.
    • If [H3O+]=5.0×10−7extM[H_3O^+] = 5.0 \times 10^{-7} ext{ M}, then:
      pH=−log⁡(5)−7=−6.3.pH = -\log(5) - 7 = -6.3.

pH of Common Substances

FluidpH
Stomach acid1.5
Lemon juice2.0
Vinegar3.0
Pure water7.0
Blood7.35 to 7.45
Household ammonia11.5

pH and pOH Relationship

  • Define pOH analogous to pH:
    • pOH=−log⁡[OH−]pOH = -\log[OH^-]
  • Relationship: pH+pOH=14pH + pOH = 14.
pH Calculations for Strong Acids and Bases

  • Strong acids/bases dissociate almost completely:

    • Example: For strong acid with concentration 0.05M, [H3O+]≈0.05extM[H_3O^+] \approx 0.05 ext{ M}; thus,
      pH=−log⁡(0.05)≈1.30.pH = -\log(0.05) \approx 1.30.

  • Practice Exercises:

    1. Calculate the $pH$ of 4.0 x 10⁻⁶ M NaOH.
    2. Calculate the $pH$ of 0.02 M HCl.

Acid-Base Equilibria

  • Reactions of weak acids/bases are treated differently due to their partial dissociation. Use ICE tables and $Ka$, $Kb$ values for calculations.

Percent Dissociation of Weak Acids

  • Percent dissociation gives a useful measure of acid strength:
    ext{% dissociation} = \frac{[HA]{ ext{dissociated}}}{[HA]{ ext{initial}}} \times 100").
  • More dilute weak acids result in higher percent dissociation due to decreased concentration of undissociated acid.

Titration Concepts

  • Titration Curve: Graphical representation of the change in $pH$ as titrant is added. Important characteristics include equivalence point where moles of acid = moles of base.

Buffers System

  • Buffers: Solutions that resist changes in $pH$ upon addition of small amounts of acid/base.
  • Henderson-Hasselbalch Equation:
    pH=pKa+log⁡([A−][HA])pH = pK_a + \log \left(\frac{[A^-]}{[HA]}\right)
  • Important for stability in biological systems.