Chapter 15 (for posting)

Chapter 15: Acids and Bases

Learning Outcomes

• Contrast Different Acid-Base Theories: Gain a comprehensive understanding of the Arrhenius, Bronsted-Lowry, and Lewis models of acids and bases to elucidate their distinct behaviors in various chemical reactions, enhancing the conceptual framework for acid-base chemistry.

• Identify Conjugate Acid-Base Pairs: Analyze and identify conjugate acid-base pairs, compare their relative strengths, and understand the intricate dynamics of proton transfer that are crucial in acid-base chemistry.

• Quantitative Comparison of Acid and Base Strengths: Compare the strengths of acids and bases using ionization constants (Ka, Kb) and/or pKa/pKb values, to quantitatively evaluate their reactivity and predict their behavior in equilibria.

• Relation of Acid Strength to Molecular Composition: Explore how structures such as electronegativity, inductive effects, resonance, and steric factors influence the ability of a molecule to donate protons, thus impacting acid strength.

• Convert Between pH and Concentration Values: Strengthen practical analytical skills by converting between pH, pOH, [H₃O⁺], and [OH⁻] for both acidic and basic solutions, which is essential for laboratory and industrial applications.

• Calculate pH and Ion Concentrations: Develop quantitative methods for assessing aqueous systems by calculating pH and ion concentrations in both strong and weak acid or base solutions, thereby improving problem-solving skills in chemistry.

• Percentage Ionization: Understand the concept of percentage ionization in weak acids to gain insights into their dissociation behavior in solution, and its practical implications for buffer systems and chemical reactions.

• pH of Salt Solutions: Determine the pH of salt solutions and analyze how the hydrolysis of constituent ions affects the overall acidity or basicity of the solution, offering deeper insights into salt behavior in various scenarios.

Acid and Base Definitions

Models of Acids and Bases:

Arrhenius Model: An early model defining:

• Acid: A substance that produces H⁺ ions in aqueous solution.Example: HCl (aq) → H⁺ (aq) + Cl⁻ (aq)

• Base: A substance that produces OH⁻ ions in aqueous solution.Example: NaOH (aq) → Na⁺ (aq) + OH⁻ (aq) Shortcomings: This model is restricted to aqueous environments and neglects the role of H⁺ ions as hydronium ions (H₃O⁺), excluding acid-base reactions occurring in non-aqueous environments.

Bronsted-Lowry Model: Broadened definition:

• Acid: Any substance that donates a proton (H⁺).Example: HCl (aq) + H₂O (l) → H₃O⁺ (aq) + Cl⁻ (aq)

• Base: A substance that accepts a proton (H⁺).Example: NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq) Key Point: Emphasizes how acids and bases function in pairs, offering insights into proton transfer mechanisms foundational to acid-base chemistry.

Lewis Model: Expands definitions to:

• Acid: Defined as an electron pair acceptor, significant for complex formation and coordination chemistry.

• Base: An electron pair donor, aiding in understanding reactions that involve Lewis acids and bases which do not fit Bronsted definitions.

Pros and Cons of Acid-Base Models

• Arrhenius Model Pros: Simple for strong acids/bases in aqueous solutions and direct relation to pH measurements.Cons: Limited to aqueous systems and does not consider organic solvents or solid states, neglecting the complexity of many real-world reactions.

• Bronsted-Lowry Model Pros: More broadly applicable, facilitating proton transfer analysis across different media, and clearly delineating donor/acceptor roles.Cons: Still limited; some reactions involving Lewis acids fall outside this model.

• Lewis Model Pros: Covers a larger variety of reactions, including those that are non-proton transfer based, and is effective in describing coordination compounds.Cons: It may be less intuitive for beginners compared to simpler definitions.

Conjugate Acid-Base Pairs

Defined as pairs of substances differing by the presence or absence of one proton. A base transforms into its conjugate acid upon proton acceptance, while an acid converts into its conjugate base when it donates a proton.Example: In the reaction between NH₃ and HCl, NH₄⁺ serves as the conjugate acid of the base NH₃, whereas Cl⁻ is the conjugate base of strong acid HCl.

Acid Strength

Strong and Weak Acids:

• Strong Acids: Completely dissociate/ionize in solution, leading to high [H₃O⁺] concentrations. Example: 1.0 M HCl provides 1.0 M [H₃O⁺] and 1.0 M [Cl⁻]. Other strong acids include HNO₃ and H₂SO₄, showing consistent behavior across concentrations.

• Weak Acids: Partially dissociate, resulting in varied [H₃O⁺] concentrations. Example: 1.0 M HF results in 0.025 M [H₃O⁺] and 0.025 M [F⁻] at equilibrium: Equilibrium reaction: HF (aq) ⇌ H⁺ (aq) + F⁻ (aq)

Measuring Acid Strength

Strength is quantitatively assessed through:

• Ka: The acid dissociation constant; a higher value indicates a stronger acid.

• pKa: The negative logarithm of Ka, wherein lower values signify stronger acids.Example Calculation: If Ka of acetic acid (CH₃COOH) is 1.8 x 10⁻⁵,

• Calculate pKa:pKa = -log(1.8 x 10⁻⁵) ≈ 4.74. Compare with HCl (pKa ≈ -7.0) to assess strength quantitatively.

Mixtures of Acids

• Strong Acid + Weak Acid: Generally, calculations skew toward strong acid behavior due to complete ionization.

• Weak Acid + Weak Acid: Evaluate using Ka values; the acid with the higher Ka predominates, significantly affecting pH.

Bases

Definitions and Strength

• Strong Bases: Fully dissociate, yielding high OH⁻ concentrations.Example: NaOH completely dissociates: NaOH → Na⁺ + OH⁻.

• Weak Bases: Partially dissociate, displaying equilibrium phenomena.Example:Equilibrium reaction: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.

Factors Influencing Base Strength

• Electron Density: Higher electron density generally enhances a base's ability to accept protons.

• Resonance Structures: Delocalization of negative charge stabilizes the base, leading to enhanced strength.

• Inductive Effects: Electronegative substituents can withdraw electron density, reducing proton acceptance efficiency.

Polyprotic Acids

Defined as substances that can donate multiple protons, with distinct ionization steps.Example: H₂SO₄ exhibits two ionization stages:

• First Ionization (Strong): H₂SO₄ → H⁺ + HSO₄⁻ (Ka1 is large)

• Second Ionization (Weak): HSO₄⁻ ⇌ H⁺ + SO₄²⁻ (Ka2 significantly smaller) Each ionization stage contributes to total [H₃O⁺], complicating acidity assessments and necessitating sequential equilibrium calculations.

pH Scale

A logarithmic scale to express solution acidity and basicity:

• pH < 7 indicates acidity.

• pH = 7 is neutral.

• pH > 7 points to basic conditions.

Calculation of pH and pOH

Using:

• pH = -log[H₃O⁺]

• pOH = -log[OH⁻]

• pH + pOH = 14 at 25°C. Example Calculation: For a 0.01 M HCl solution:

• Determine [H₃O⁺]: HCl is a strong acid, yielding [H₃O⁺] = 0.01 M.

• Calculate pH: pH = -log(0.01) = 2.

Water's Role

Water exhibits amphoteric behavior, acting as both an acid and a base, crucial for numerous acid-base reactions. Its autoionization significantly contributes to establishing [H₃O+] and [OH⁻]: Reaction: H₂O ⇌ H₃O⁺ + OH⁻.

Salt Solutions

Salt solutions can exhibit neutral, acidic, or basic properties based on their constituent ions:

• Neutral Salt: NaCl remains neutral without significant reactions with water.

• Acidic Salt: NH₄Cl becomes acidic via NH₄⁺ hydrolysis.

• Basic Salt: NaCH₃CO₂ leads to basicity through CH₃COO⁻ hydrolysis. Calculating pH in salt solutions involves assessing hydrolysis dynamics: Example: Determine pH of a 0.1 M NH₄Cl solution:

• Identify [H₃O⁺] due to hydrolysis: NH₄⁺ + H₂O ⇌ H₃O⁺ + NH₃.

• Determine Kb for NH₃ to find Ka for NH₄⁺ (using Kw = 1.0 x 10⁻¹⁴).

• Implement ICE table methods to resolve [H₃O⁺] and derive pH.

Practice Questions

1. Strength Comparison:Compare the strength of the following acids and explain your reasoning:

   • HCl

   • HF

2. Calculate pH:Calculate the pH of a 0.005 M HCl solution. Show your calculations.

3. Percent Ionization:If the ionization constant (Ka) for acetic acid is 1.8 x 10⁻⁵, calculate the percentage of ionization if 0.1 M acetic acid ionizes to produce [H₃O⁺].

4. Effect of Structure:How does the presence of electronegative atoms in a molecule affect its acid strength? Provide an example.

5. Salt Solution pH:Determine the pH of a 0.1 M solution of NH₄Cl. Explain the steps you would take to arrive at your answer.

6. Arrhenius vs. Bronsted-Lowry:Describe the key differences between the Arrhenius and Bronsted-Lowry definitions of acids and bases.

7. Ka and pKa:Given that the Ka of hydrofluoric acid (HF) is 6.8 x 10⁻⁴, calculate its pKa.

8. Polyprotic Acids:Explain the differences in the ionization steps of a diprotic acid versus a monoprotic acid, using H₂SO₄ as an example.

9. Acid-Base Strength in Solutions:If you mixed equal concentrations of HCl and acetic acid, predict the resulting pH of the solution and justify your reasoning based on the strength of the acids.

10. Buffer System Example:Describe how a buffer system works using a specified weak acid and its conjugate base. Include a specific example, such as acetic acid and sodium acetate, and explain how it regulates pH changes.

11. Real-World Application:Explain how acid-base reactions are utilized in biological systems, such as the bicarbonate buffer system in blood. Discuss the balance of pH and its significance to human health.

12. Le Chatelier's Principle:How would adding a strong acid affect the equilibrium position of the dissociation of a weak acid in solution? Use Le Chatelier's Principle to explain your answer.

13. Electronegativity and Acid Strength:Choose a series of acids (e.g., HF, HCl, HBr, HI) and analyze how the increasing size and decreasing electronegativity of the halogen atom affects the acid strength. Include a detailed explanation of the trend in H-X bond strength.

14. Calculating pH of Salt Solutions:A 0.1 M solution of sodium acetate is prepared. Calculate the pH, knowing that the Kb of acetate ion (CH₃COO⁻) is 5.6 x 10⁻¹¹. Show all steps in your calculation.

15. Polyprotic Acid Behavior:Discuss how the multiple ionization constants for polyprotic acids affect their behavior in solutions. Compare the first and second dissociation constants of sulfuric acid (H₂SO₄) and how they contribute to overall acidity.

16. Acid-Base Titration Concepts:In a titration of 0.1 M acetic acid with 0.1 M NaOH, describe the changes in pH that occur throughout the titration process. Include the concept of the equivalence point and how it relates to acid-base strength.

17. Ionization and Solvent Effects:How does the choice of solvent influence the ionization of acids and bases? Discuss the differences in behavior for strong acids in non-aqueous solvents compared to aqueous solutions.

18. Acid-Base Reactions in Industry:Identify a specific industrial process that involves acid-base reactions (such as the manufacturing of fertilizers or pharmaceuticals) and discuss the role of acid-base chemistry in that process.