ACID-BASE EQUILIBRIA

Chapter 14 ACID-BASE EQUILIBRIA

14.1 Brønsted-Lowry Acids and Bases

• Definition of acids and bases: According to the Brønsted-Lowry acid-base theory:

  • An acid is a substance that donates a proton (H+).

  • A base is a substance that accepts a proton (H+).

• Learning Objective: Identify acids, bases, and conjugate acid-base pairs according to the Brønsted-Lowry definition.

• Ionization Reactions: Write equations for acid and base ionization reactions.

  • Example:** For hydrofluoric acid (HF):
    $HF(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + F^-(aq)$

• Calculating Hydronium and Hydroxide Ion Concentrations: Use the ion-product constant for water ($K_w$) to calculate hydronium and hydroxide ion concentrations.

  • $K<em>w = [H</em>3O^+][OH^-] = 1.0 imes 10^{-14}$ at 25°C

• Amphiprotic Substances: Describe the acid-base behavior of amphiprotic substances, which can act as either an acid or a base based on the reaction context.

14.2 pH and pOH

• Learning Objective: Explain the characterization of aqueous solutions as acidic, basic, or neutral and express hydronium and hydroxide ion concentrations on the pH and pOH scales.

• The pH Scale:

  • Used to describe the acidity of solutions with values typically ranging from 0 to 14.

  • A pH of less than 7 indicates acidity, a pH of 7 is neutral, and a pH greater than 7 indicates basicity.

• Measuring pH: pH can be determined using:

  • A pH meter.

  • pH paper or indicators that change color at specific pH values.

• Calculating pH:

  • The relationship is defined by the equation:
    $pH = -\log[H_3O^+]$

  • Example: If $[H_3O^+] = 1.0 \times 10^{-2}$ M, then $pH = -\log(1.0 \times 10^{-2}) = 2.00$.

• Calculating [H_3O+] from pH: For whole number pH values,

  • $[H_3O^+] = 10^{-pH}$.

• Significant Figures in pH:

  • The number of decimal places in the pH value corresponds to the significant figures of [H3O+].

14.3 Relative Strengths of Acids and Bases

• Strong vs. Weak Acids:

  • A strong acid completely ionizes in solution:

  • Example: $HCl(g) + H<em>2O(l) \rightarrow H</em>3O^+(aq) + Cl^-(aq)$

  • A weak acid only partially ionizes:

  • Example: $H<em>2CO</em>3(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + HCO_3^-(aq)$

• Weak Acids: Weak acids have strong conjugate bases and do not dissociate completely in water, generally resulting in lower concentrations of H3O+ and the anion.

• Characteristics of Bases:

  • Arrhenius bases produce hydroxide ions (OH-) in water.

  • Common examples of strong bases include:

  • $NaOH$ (sodium hydroxide), $KOH$ (potassium hydroxide), and $Ca(OH)_2$ (calcium hydroxide).

• Naming Acids:

  • Acids containing hydrogen and a nonmetal are named with the prefix “hydro” and end with “ic”:

  • Example: HCl is called hydrochloric acid.

  • Acids with a hydrogen ion and a polyatomic ion derived from the -ate ion change to -ic and from the -ite ion change to -ous.

Learning Check Solutions

• Identifying Characteristics of Acids and Bases:

  • A: Acid - has a sour taste.

  • B: Base - produces OH- in aqueous solutions.

  • C: Base - has a chalky taste.

  • D: Acid & Base - both are electrolytes.

  • E: Acid - produces H+ in aqueous solutions.

• Conjugate Acid-Base Pairs: In any acid-base reaction, there are two conjugate acid-base pairs related by the loss and gain of H+:

  • Example of pairing with identification:

  • In the reaction $HNO<em>3(aq) + NH</em>3(aq) \rightarrow NO<em>3^−(aq) + NH</em>4^+(aq)$,

    • $HNO3$ is the acid, and $NO3^-$ is its conjugate base; $NH3$ is the base, and $NH4^+$ is its conjugate acid.