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GOB Chemistry: Acid-Base Equilibria and Buffers

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Termini in questo insieme (20)

  • Brønsted-Lowry acid

    A proton donor in a chemical reaction.

  • Brønsted-Lowry base

    A proton acceptor in a chemical reaction.

  • Conjugate acid-base pair

    Two species differing by one proton; the acid donates a proton to form its conjugate base, and the base accepts a proton to form its conjugate acid.

  • Auto-ionization of water

    Water molecules react with each other to form \(\mathrm{H_3O^+}\) and \(\mathrm{OH^-}\), establishing an equilibrium.

  • Ion-product constant of water, Kw

    At 25°C, \(K_w = [H^+][OH^-] = 1.0 \times 10^{-14}\).

  • pH definition

    \(\mathrm{pH} = -\log[H^+] = -\log[H_3O^+]\), a measure of hydrogen ion concentration.

  • Relationship between pH and pOH

    \(\mathrm{pH} + \mathrm{pOH} = 14.0\) at 25°C.

  • Strong acid characteristics

    Completely ionizes in water, strong electrolytes, large equilibrium constants, e.g., HCl, HNO3, H2SO4.

  • Weak acid characteristics

    Partially ionizes in water, weak electrolytes, smaller equilibrium constants, e.g., acetic acid (CH3COOH).

  • Acid dissociation constant, Ka

    \(K_a = \frac{[H_3O^+][A^-]}{[HA]}\), measures acid strength; larger Ka means stronger acid.

  • Polyprotic acids

    Acids that can donate more than one proton, e.g., diprotic (H2SO4) and triprotic (H3PO4) acids.

  • Base dissociation constant, Kb

    Equilibrium constant for base ionization; larger Kb means stronger base.

  • Relationship between Ka and Kb

    \(K_a \times K_b = K_w = 1.0 \times 10^{-14}\) at 25°C for conjugate acid-base pairs.

  • Amphiprotic substances

    Can act as either an acid or a base, e.g., water (H2O).

  • Common ion effect

    Presence of a common ion suppresses ionization of a weak electrolyte, shifting equilibrium according to Le Chatelier's principle.

  • Buffer solution

    A solution containing a weak acid and its conjugate base (or weak base and its conjugate acid) that resists pH changes upon addition of small amounts of acid or base.

  • Henderson-Hasselbalch equation

    \(\mathrm{pH} = pK_a + \log \frac{[A^-]}{[HA]}\), used to calculate pH of buffer solutions.

  • Buffer capacity

    The amount of acid or base a buffer can neutralize before pH changes appreciably; depends on concentrations of acid and conjugate base.

  • Effect of adding strong acid to acetate buffer

    Equilibrium shifts left, acetate ions react with added H+, pH remains nearly constant.

  • Effect of adding strong base to acetate buffer

    OH- reacts with H3O+, equilibrium shifts right, acetic acid dissociates to restore H3O+, pH remains nearly constant.