Acids and Bases

Acid-Base Calculations

  1. pH and pOH

    • pH is a measure of the concentration of hydrogen ions ([H+][H^+]) in a solution.

    • pOH is a measure of the concentration of hydroxide ions ([OH−][OH^-]) in a solution.

    • pH=−log10[H+]pH = -log_{10}[H^+]

    • pOH=−log10[OH−]pOH = -log_{10}[OH^-]

    • In aqueous solutions at 25∘C25^\circ C, pH+pOH=14pH + pOH = 14

  2. Strong Acids and Bases

    • Strong acids and bases completely dissociate in water.

    • For strong acids, the concentration of H+H^+ is equal to the concentration of the acid.

    • For strong bases, the concentration of OH−OH^- is equal to the concentration of the base.

      • Example: If you have a 0.01 M solution of HClHCl, then [H+]=0.01M[H^+] = 0.01 M and pH=−log10(0.01)=2pH = -log_{10}(0.01) = 2

  3. Weak Acids and Bases

    • Weak acids and bases only partially dissociate in water.

    • The acid dissociation constant, KaK_a, is used to measure the strength of a weak acid.

    • The base dissociation constant, KbK_b, is used to measure the strength of a weak base.

    • Ka=[H+][A−][HA]K_a = \frac{[H^+][A^-]}{[HA]}, where HA is the weak acid and A- is its conjugate base.

    • Kb=[OH−][HB+][B]K_b = \frac{[OH^-][HB^+]}{[B]}, where B is the weak base and HB+ is its conjugate acid.

    • The smaller the K<em>aK<em>a or K</em>bK</em>b, the weaker the acid or base.

  4. Calculations with Weak Acids and Bases

    • To calculate the pH of a weak acid or base solution, you typically need to use an ICE table.

    • ICE (Initial, Change, Equilibrium) table helps to organize the concentrations of the species in the equilibrium.

    • Example: Calculate the pH of a 0.1 M solution of acetic acid (CH<em>3COOHCH<em>3COOH, K</em>a=1.8×10−5K</em>a = 1.8 \times 10^{-5})

      • ICE Table:

        CH3COOHCH_3COOH

        H+H^+

        CH3COO−CH_3COO^-

        Initial (I)

        0.1

        0

        0

        Change (C)

        -x

        +x

        +x

        Equilib (E)

        0.1 - x

        x

        x

      • K<em>a=[H+][CH</em>3COO−][CH3COOH]=x20.1−x=1.8×10−5K<em>a = \frac{[H^+][CH</em>3COO^-]}{[CH_3COOH]} = \frac{x^2}{0.1 - x} = 1.8 \times 10^{-5}

      • Assume x is small, so 0.1 - x ≈ 0.1

      • x2=1.8×10−6x^2 = 1.8 \times 10^{-6}

      • x=1.8×10−6=0.00134M=[H+]x = \sqrt{1.8 \times 10^{-6}} = 0.00134 M = [H^+]

      • pH=−log10(0.00134)=2.87pH = -log_{10}(0.00134) = 2.87

  5. Acid-Base Titrations

    • Titration is a process used to determine the concentration of an acid or base by neutralizing it with a known concentration of a base or acid.

    • Equivalence point: The point at which the acid and base have completely reacted with each other.

    • Endpoint: The point at which the indicator changes color.

    • For a strong acid-strong base titration, the pH at the equivalence point is 7.

    • For a weak acid-strong base titration, the pH at the equivalence point is greater than 7.

    • For a strong acid-weak base titration, the pH at the equivalence point is less than 7.

  6. Buffers

    • A buffer is a solution that resists changes in pH when small amounts of acid or base are added.

    • A buffer typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid.

    • The pH of a buffer can be calculated using the Henderson-Hasselbalch equation:

      • pH=pK<em>a+log</em>10([A−][HA])pH = pK<em>a + log</em>{10}(\frac{[A^-]}{[HA]})

      • pOH=pK<em>b+log</em>10([HB+][B])pOH = pK<em>b + log</em>{10}(\frac{[HB^+]}{[B]})

      • where pK<em>a=−log</em>10(K<em>a)pK<em>a = -log</em>{10}(K<em>a) and pK</em>b=−log<em>10(K</em>b)pK</em>b = -log<em>{10}(K</em>b)

  7. Polyprotic Acids

    • Acids that have more than one ionizable proton are called polyprotic acids (e.g