pH Practice

  1. Acid = gives away protons (H+H^+)

    • Arrhenius definition: Releases H+H^+ ions in aqueous solution to form hydronium ions (H3O+H_3O^+).

    • Brønsted-Lowry definition: Proton donor (H+H^+ donor).

    • Strong acids (e.g., HCl(aq)HCl_{(aq)}, HNO3(aq)HNO_{3(aq)}) dissociate completely; weak acids (e.g., CH3COOH(aq)CH_3COOH_{(aq)}) dissociate partially.

  2. Base = accepts protons (H+H^+)

    • Arrhenius definition: Releases hydroxide ions (OH−OH^-) in aqueous solution.

    • Brønsted-Lowry definition: Proton acceptor (H+H^+ acceptor).

    • Alkalis are soluble bases that produce OH−OH^- ions in water (e.g., NaOH(aq)NaOH_{(aq)}, KOH(aq)KOH_{(aq)}).

  3. More H+H^+ → pH goes DOWN

    • pH is a logarithmic measure of hydrogen ion concentration: pH=−log10([H+])\text{pH} = -\text{log}_{10}([H^+]).

    • Each single unit change on the pH scale represents a 10-fold change in H+H^+ concentration (e.g., pH 3 has 10 times more H+H^+ than pH 4).

  4. Less H+H^+ → pH goes UP

    • As H+H^+ concentration decreases, the solution becomes less acidic and more basic/alkaline.

And for hydroxide:

  • More OH−OH^- → pH goes UP

    • High OH−OH^- concentration correlates to low H+H^+ concentration, increasing pH toward 14.

  • Less OH−OH^- → pH goes DOWN

    • Decreasing OH−OH^- concentration increases relative H+H^+ concentration, decreasing pH toward 0.

    • The ionic product of water maintains this equilibrium: Kw=[H+][OH−]=1.0×10−14 mol2 dm−6K_w = [H^+][OH^-] = 1.0 \times 10^{-14}\,mol^2\,dm^{-6} at 298 K298\,K.

pH scale

0 ← acidic ——— 7 neutral ——— basic → 14

  • Toward 0 = high H+H^+, low OH−OH^- (strongly acidic, e.g., battery acid, gastric juice).

  • Toward 14 = low H+H^+, high OH−OH^- (strongly basic/alkaline, e.g., drain cleaner, bleach).

  • At 7 = H+H^+ and OH−OH^- concentrations are equal ([H+]=[OH−]=1.0×10−7 mol dm−3[H^+] = [OH^-] = 1.0 \times 10^{-7}\,mol\,dm^{-3} at 298 K298\,K).