Acids, Bases, and Equilibrium Notes

Acids and Bases

  • The term "acid" comes from the Latin word acidus, meaning “sour”.
  • Arrhenius acids:
    • Produce hydrogen ions (H+H^+) when dissolved in water.
    • Are electrolytes.
    • Turn blue litmus red.
    • Corrode some metals.
  • Arrhenius bases:
    • Produce hydroxide ions (OH−OH^-) in water.
    • Taste bitter or chalky.
    • Are electrolytes.
    • Feel soapy and slippery.
    • Turn litmus paper blue and phenolphthalein pink.

Naming Acids

  • Acids with hydrogen (H+H^+) and a nonmetal (or CN−CN^-) use the prefix hydro and end with -ic acid (e.g., HCl hydrochloric acid).
  • Acids with hydrogen (H+H^+) and a polyatomic ion:
    • -ate changes to -ic acid.
    • -ite changes to -ous acid.
    • Example: ClO<em>3−ClO<em>3^- (chlorate) becomes HClO</em>3HClO</em>3 (chloric acid); ClO<em>2−ClO<em>2^- (chlorite) becomes HClO</em>2HClO</em>2 (chlorous acid).

Naming Bases

  • Arrhenius bases are named as hydroxides (e.g., NaOH sodium hydroxide).

Brønsted–Lowry Acids and Bases

  • Acid: donates H+H^+
  • Base: accepts H+H^+

Conjugate Acid–Base Pairs

  • Two pairs in any acid-base reaction.
  • Each pair is related by the loss and gain of H+H^+.
  • One pair occurs in the forward direction, one in the reverse.
  • Example: NH<em>3/NH</em>4+NH<em>3/NH</em>4^+ and H2O/OH−H_2O/OH^-.

Strengths of Acids and Bases

  • Strong acids: completely ionize in water.
  • Weak acids: partially dissociate in water.
  • Strong bases: formed from Group 1A and 2A metals, dissociate completely in water.
  • Weak bases: weak electrolytes, produce few ions in solution (e.g., NH3NH_3).

Equilibrium

  • Reversible reaction proceeds in both forward and reverse directions.
  • Equilibrium is reached when there are no further changes in concentrations of reactants/products, and the rate of forward reaction equals the rate of reverse reaction.

Le Châtelier’s Principle

  • When equilibrium is disturbed, the rates of forward and reverse reactions change to relieve the stress and reestablish equilibrium.
  • Adding a substance shifts the reaction away from that substance.

Dissociation of Water

  • Water is amphoteric (can act as acid or base).
  • H<em>2O(l)+H</em>2O(l)&amp;leftrightarrowH3O+(aq)+OH−(aq)H<em>2O(l) + H</em>2O(l) \&amp;leftrightarrow H_3O^+(aq) + OH^-(aq)
  • In pure water at 25°C: [H3O+]=[OH−]=1.0x10−7M[H_3O^+] = [OH^-] = 1.0 x 10^{-7} M
  • K<em>w=[H</em>3O+][OH−]=1.0x10−14K<em>w = [H</em>3O^+][OH^-] = 1.0 x 10^{-14}
  • Neutral solution: [H3O+]=[OH−][H_3O^+] = [OH^-]
  • Acidic solution: [H_3O^+] > [OH^-]
  • Basic solution: [OH^-] > [H_3O^+]
  • [OH−]=K<em>w/[H</em>3O+][OH^-] = K<em>w / [H</em>3O^+] and [H<em>3O+]=K</em>w/[OH−][H<em>3O^+] = K</em>w / [OH^-]

The pH Scale

  • Indicates acidity of a solution, usually ranges from 0 to 14.
  • pH = −log[H3O+]−log[H_3O^+]
  • Acidic: pH < 7
  • Neutral: pH = 7
  • Basic: pH > 7
  • A change of one pH unit corresponds to a tenfold change in [H3O+][H_3O^+].
  • [H3O+]=10−pH[H_3O^+] = 10^{-pH}

Reactions of Acids

  • With certain metals to produce a salt and hydrogen gas.
  • With bases to produce a salt and water (neutralization).
  • With bicarbonate and carbonate ions to produce carbon dioxide gas.

Neutralization

  • Acid + Base → Salt + Water
  • H+(aq)+OH−(aq)→H2O(l)H^+(aq) + OH^-(aq) → H_2O(l)

Acid–Base Titration

  • Used to determine the molarity of an acid.
  • Uses a base (e.g., NaOH) to neutralize a measured volume of acid.
  • Endpoint: moles of base = moles of acid.

Buffers

  • Maintain pH by neutralizing small amounts of added acid or base.
  • Contain a weak acid and a salt of its conjugate base.
  • Resist changes in pH.
  • Weak acid neutralizes bases; conjugate base neutralizes acids.