Acids and Bases Chapter Review

Learning Outcomes

• Contrast the Arrhenius, Bronsted-Lowry, and Lewis models of acids/bases.

• Identify conjugate acid-base pairs and compare their relative strengths.

• Compare strengths of acids and bases through their ionization constants (Ka, Kb) or pKa/pKb values.

• Relate the strength of acids to their molecular composition.

• Convert between pH, pOH, [H3O+], [OH-] for acidic and basic solutions.

• Calculate the pH and/or ion concentrations of solutions containing a strong acid, strong base, weak acid, or weak base.

• Calculate the % ionization of weak acids.

• Calculate the pH and/or ion concentrations of acid mixtures.

• Calculate the pH and/or ion concentrations of polyprotic acids.

• Determine the pH of a salt solution.

Models of Acids and Bases

Arrhenius Definition

• Acid: A substance that produces H+ ions in aqueous solution.

  • Example: $\text{HCl (aq) } \rightarrow \text{ H}^+ (aq) + \text{ Cl}^- (aq)$

• Base: A substance that produces OH- ions in aqueous solution.

  • Example: $\text{NaOH (aq)} \rightarrow \text{Na}^+ (aq) + \text{OH}^- (aq)$

• Limitations:

  • Does not account for H+ existing as H3O+ in water.

  • Limited to aqueous solutions.

Bronsted-Lowry Definition

• Acid: A substance that donates a proton (H+).

  • Example: $\text{HCl (aq) + H2O (l) } \rightarrow \text{ H3O}^+ (aq) + \text{ Cl}^- (aq)$

• Base: A substance that accepts a proton (H+).

  • Example: $\text{NH}<em>3 (aq) + \text{H2O (l)} \rightleftharpoons \text{NH}</em>4^+ (aq) + \text{OH}^- (aq)$

• Collaboration: Acids and bases always work together in the Bronsted-Lowry model.

Lewis Definition

• Lewis Acid: An electron pair acceptor.

• Lewis Base: An electron pair donor.

• Expands the definition to include reactions that do not involve proton transfers.

Conjugate Acid-Base Pairs

• Definition: Two substances differing only by the presence/absence of a proton.

• A base accepts a proton to become a conjugate acid.

• An acid donates a proton to become a conjugate base.

Acid Strength

• Strong Acids: Completely dissociate in solution.

  • Example: $1.0 \text{ M HCl} \rightarrow 1.0 \text{ M [H3O+]} \, ext{and} \, 1.0 \text{ M [Cl-]}$

• Weak Acids: Only partially dissociate in solution.

  • Example: $1.0 \text{ M HF} \rightarrow 0.025 \text{ M [H3O+]} \, ext{and} \, 0.025 \text{ M [F-]}$

• Dissociation and Ionization Constants:

  • Strong acids have larger Ka, while weak acids have smaller Ka.

Measuring Acid Strength

• General Rule: Strong acids yield relatively large Ka values; weak acids yield small Ka values.

• Relationship between acid strength and conjugate base strength:

  • As strength of HA increases, its conjugate base A- becomes weaker.

Molecular Structure Influences on Acidity

• Binary Acids: Factors include bond strength and stability of conjugate base.

  • Example comparative values:

  • \text{H-F} < \text{H-Cl} < \text{H-Br} < \text{H-I}

• Oxoacids: Strength depends on electronegativity of the central atom and number of O atoms affecting the O-H bond.

• More electronegative central atoms weaken the O-H bond, increasing acidity.

Calculating pH

• Strong Acids and Bases: Use direct concentration to find pH. Strong acids are straightforward as they fully dissociate.

• Weak Acids: Use ICE tables and associated Ka values to determine pH.

  • For example: For $\text{HA} (aq)$ in equilibrium with $\text{H3O+} (aq)$ and $\text{A-} (aq)$, set up equilibrium expressions and solve for x to compute pH.

Percent Ionization of Weak Acids

• Expressed as:
  \text{% Ionization} = \frac{[\text{Ionized Acid}]}{[\text{Acid}]_{initial}} \times 100%

• Increases with increasing Ka and decreasing initial concentration.

Mixtures of Acids

1. Strong + Weak Acid: The strong acid primarily determines the pH.

2. Weak + Weak Acid: The strongest weak acid significantly contributes to [H3O+]. Use equilibrium calculations for analysis.

pH of Salt Solutions

• The pH is determined by the nature of the ions derived from the salt:

  • Neutral salts: do not alter pH (e.g., NaCl).

  • Acidic salts: produce acidic solutions (e.g., NH4Cl).

  • Basic salts: produce basic solutions (e.g., NaCH3CO2).

  • Hydrolysis reactions govern the pH behavior of anions and cations.

To identify the type of salt, you should consider the ions that make up the salt and their origins:

1. Neutral Salts: Formed from the reaction of a strong acid and a strong base. They do not affect the pH of the solution (e.g., NaCl).

2. Acidic Salts: Formed from a strong acid and a weak base. These salts produce acidic solutions when dissolved in water (e.g., NH4Cl).

3. Basic Salts: Formed from a weak acid and a strong base. They produce basic solutions when dissolved in water (e.g., NaCH3CO2).

The pH behavior of salts is influenced by hydrolysis reactions involving their anions and cations.

Practice Problems and Applications

• Practice questions are included throughout the content.

• When performing calculations, carefully track initial concentrations, changes, and set up appropriate equilibrium expressions.