Chemistry Notes: Acids and Bases

Chapter 14: Acids and Bases

14.1 Nature of Acids and Bases

  • Acids & Bases: Different definitions

    • Arrhenius: Acids produce H^+ ions; bases produce OH^- ions.

    • Brønsted-Lowry: Acids are proton (H+) donors; bases are proton acceptors.

    • Lewis: Acids are electron acceptors; bases are electron donors.

14.2 pH Scale

  • pH Definition: pH = - ext{log}[H^+]

  • pOH: pOH = - ext{log}[OH^-] and pH + pOH = 14.00

  • Acidic/Basic Nature:

    • pH < 7: acidic

    • pH = 7: neutral

    • pH > 7: basic

14.3 Calculating pH

  • Strong Acids: Fully dissociate in solution (i.e., HCl).

  • Weak Acids: Partially dissociate; ionization results in equilibrium.

  • Acid-Dissociation Constant: K_a = \frac{[H^+][A^-]}{[HA]}

14.4 Properties of Bases

  • Strong Bases: Fully dissociating bases, such as NaOH.

  • Weak Bases: Partially dissociating bases, like NH_3.

14.5 Acid-Base Properties of Salts

  • Neutral Solutions: Salts from strong acids and bases (e.g., NaCl).

  • Basic Solutions: When anions are conjugate bases of weak acids (e.g., NaF).

  • Acidic Solutions: When cations are conjugate acids of weak bases (e.g., NH_4Cl).

14.6 Additional Acid-Base Properties

  • Acidity increases with higher electronegativity and lower bond strength in nonmetal hydrides.

14.7 Lewis Acid-Base Model

  • Acids accept electrons; bases donate electrons.

14.8 Strategies for Solving Problems

  • Identify major species and reactions.

  • Use appropriate equations for equilibria.

  • Determine concentration changes and solve for [H^+] and pH.

14.1 Nature of Acids and Bases
Acids & Bases: Different definitions
Arrhenius: Acids produce H^+ ions; bases produce OH^- ions.
Brønsted-Lowry: Acids are proton (H+) donors; bases are proton acceptors.
Lewis: Acids are electron acceptors; bases are electron donors.
Example:

  • Hydrochloric acid (HCl) is an Arrhenius acid because it produces H^+ ions in water, while sodium hydroxide (NaOH) is an Arrhenius base because it produces OH^- ions.

14.2 pH Scale
pH Definition: pH = - \text{log}[H^+]
pOH: pOH = - \text{log}[OH^-] and pH + pOH = 14.00
Acidic/Basic Nature: pH < 7: acidic pH = 7: neutral pH > 7: basic
Example:

  • A solution with a pH of 3 is acidic, while a solution with a pH of 9 is basic.

14.3 Calculating pH
Strong Acids: Fully dissociate in solution (i.e., HCl).
Weak Acids: Partially dissociate; ionization results in equilibrium.
Acid-Dissociation Constant: K_a = \frac{[H^+][A^-]}{[HA]}
Example:

  • For a 0.01 M HCl solution, since it is a strong acid, pH = 2. For acetic acid, the K_a allows for equilibrium concentration calculations.

14.4 Properties of Bases
Strong Bases: Fully dissociating bases, such as NaOH.
Weak Bases: Partially dissociating bases, like NH_3.
Example:

  • Sodium hydroxide (NaOH) completely dissociates in water, whereas ammonia (NH₃) only partially dissociates, resulting in a less basic solution.

14.5 Acid-Base Properties of Salts
Neutral Solutions: Salts from strong acids and bases (e.g., NaCl).
Basic Solutions: When anions are conjugate bases of weak acids (e.g., NaF).
Acidic Solutions: When cations are conjugate acids of weak bases (e.g., NH_4Cl).
Example:

  • The reaction of $NaCl$ in water results in a neutral solution, while $NH4Cl$ creates an acidic solution due to the presence of $NH4^+$ ions.

14.6 Additional Acid-Base Properties
Acidity increases with higher electronegativity and lower bond strength in nonmetal hydrides.
Example:

  • Hydrofluoric acid (HF) is a stronger acid than acetic acid (CH₃COOH) due to the electronegativity of fluorine.

14.7 Lewis Acid-Base Model
Acids accept electrons; bases donate electrons.
Example:

  • The reaction of boron trifluoride (BF₃) with ammonia (NH₃) demonstrates Lewis acid-base behavior; BF₃ is the Lewis acid as it accepts an electron pair from NH₃, the Lewis base.

14.8 Strategies for Solving Problems
Identify major species and reactions.
Use appropriate equations for equilibria.
Determine concentration changes and solve for [H^+] and pH.
Example:

  • In a titration problem, to find the final pH, identify the weak acid and strong base involved, then apply the corresponding equilibrium equations to compute the resulting pH after neutralization.

14.1 Nature of Acids and Bases
Acids & Bases: Different definitions
Arrhenius: Acids produce H^+ ions; bases produce OH^- ions.
Brønsted-Lowry: Acids are proton (H+) donors; bases are proton acceptors.
Lewis: Acids are electron acceptors; bases are electron donors.
Example:

  • Hydrochloric acid (HCl) is an Arrhenius acid because it produces H^+ ions in water, while sodium hydroxide (NaOH) is an Arrhenius base because it produces OH^- ions.

14.2 pH Scale
pH Definition: pH = - \text{log}[H^+]
pOH: pOH = - \text{log}[OH^-] and pH + pOH = 14.00
Acidic/Basic Nature: pH < 7: acidic pH = 7: neutral pH > 7: basic
Example:

  • A solution with a pH of 3 is acidic, while a solution with a pH of 9 is basic.

14.3 Calculating pH
Strong Acids: Fully dissociate in solution (i.e., HCl).
Weak Acids: Partially dissociate; ionization results in equilibrium.
Acid-Dissociation Constant: K_a = \frac{[H^+][A^-]}{[HA]}
Example:

  • For a 0.01 M HCl solution, since it is a strong acid, pH = 2. For acetic acid, the K_a allows for equilibrium concentration calculations.

14.4 Properties of Bases
Strong Bases: Fully dissociating bases, such as NaOH.
Weak Bases: Partially dissociating bases, like NH_3.
Example:

  • Sodium hydroxide (NaOH) completely dissociates in water, whereas ammonia (NH₃) only partially dissociates, resulting in a less basic solution.

14.5 Acid-Base Properties of Salts
Neutral Solutions: Salts from strong acids and bases (e.g., NaCl).
Basic Solutions: When anions are conjugate bases of weak acids (e.g., NaF).
Acidic Solutions: When cations are conjugate acids of weak bases (e.g., NH_4Cl).
Example:

  • The reaction of $NaCl$ in water results in a neutral solution, while $NH4Cl$ creates an acidic solution due to the presence of $NH4^+$ ions.

14.6 Additional Acid-Base Properties
Acidity increases with higher electronegativity and lower bond strength in nonmetal hydrides.
Example:

  • Hydrofluoric acid (HF) is a stronger acid than acetic acid (CH₃COOH) due to the electronegativity of fluorine.

14.7 Lewis Acid-Base Model
Acids accept electrons; bases donate electrons.
Example:

  • The reaction of boron trifluoride (BF₃) with ammonia (NH₃) demonstrates Lewis acid-base behavior; BF₃ is the Lewis acid as it accepts an electron pair from NH₃, the Lewis base.

14.8 Strategies for Solving Problems
Identify major species and reactions.
Use appropriate equations for equilibria.
Determine concentration changes and solve for [H^+] and pH.
Example:

  • In a titration problem, to find the final pH, identify the weak acid and strong base involved, then apply the corresponding equilibrium equations to compute the resulting pH after neutralization.

14.1 Nature of Acids and Bases
Acids & Bases: Different definitions
Arrhenius: Acids produce H^+ ions; bases produce OH^- ions.
Brønsted-Lowry: Acids are proton (H+) donors; bases are proton acceptors.
Lewis: Acids are electron acceptors; bases are electron donors.
Example:

  • Hydrochloric acid (HCl) is an Arrhenius acid because it produces H^+ ions in water, while sodium hydroxide (NaOH) is an Arrhenius base because it produces OH^- ions.

14.2 pH Scale
pH Definition: pH = - \text{log}[H^+]
pOH: pOH = - \text{log}[OH^-] and pH + pOH = 14.00
Acidic/Basic Nature: pH < 7: acidic pH = 7: neutral pH > 7: basic
Example:

  • A solution with a pH of 3 is acidic, while a solution with a pH of 9 is basic.

14.3 Calculating pH
Strong Acids: Fully dissociate in solution (i.e., HCl).
Weak Acids: Partially dissociate; ionization results in equilibrium.
Acid-Dissociation Constant: K_a = \frac{[H^+][A^-]}{[HA]}
Example:

  • For a 0.01 M HCl solution, since it is a strong acid, pH = 2. For acetic acid, the K_a allows for equilibrium concentration calculations.

14.4 Properties of Bases
Strong Bases: Fully dissociating bases, such as NaOH.
Weak Bases: Partially dissociating bases, like NH_3.
Example:

  • Sodium hydroxide (NaOH) completely dissociates in water, whereas ammonia (NH₃) only partially dissociates, resulting in a less basic solution.

14.5 Acid-Base Properties of Salts
Neutral Solutions: Salts from strong acids and bases (e.g., NaCl).
Basic Solutions: When anions are conjugate bases of weak acids (e.g., NaF).
Acidic Solutions: When cations are conjugate acids of weak bases (e.g., NH_4Cl).
Example:

  • The reaction of $NaCl$ in water results in a neutral solution, while $NH4Cl$ creates an acidic solution due to the presence of $NH4^+$ ions.

14.6 Additional Acid-Base Properties
Acidity increases with higher electronegativity and lower bond strength in nonmetal hydrides.
Example:

  • Hydrofluoric acid (HF) is a stronger acid than acetic acid (CH₃COOH) due to the electronegativity of fluorine.

14.7 Lewis Acid-Base Model
Acids accept electrons; bases donate electrons.
Example:

  • The reaction of boron trifluoride (BF₃) with ammonia (NH₃) demonstrates Lewis acid-base behavior; BF₃ is the Lewis acid as it accepts an electron pair from NH₃, the Lewis base.

14.8 Strategies for Solving Problems
Identify major species and reactions.
Use appropriate equations for equilibria.
Determine concentration changes and solve for [H^+] and pH.
Example:

  • In a titration problem, to find the final pH, identify the weak acid and strong base involved, then apply the corresponding equilibrium equations to compute the resulting pH after neutralization.

14.1 Nature of Acids and Bases
Acids & Bases: Different definitions
Arrhenius: Acids produce H^+ ions; bases produce OH^- ions.
Brønsted-Lowry: Acids are proton (H+) donors; bases are proton acceptors.
Lewis: Acids are electron acceptors; bases are electron donors.
Example:

  • Hydrochloric acid (HCl) is an Arrhenius acid because it produces H^+ ions in water, while sodium hydroxide (NaOH) is an Arrhenius base because it produces OH^- ions.

14.2 pH Scale
pH Definition: pH = - \text{log}[H^+]
pOH: pOH = - \text{log}[OH^-] and pH + pOH = 14.00
Acidic/Basic Nature: pH < 7: acidic pH = 7: neutral pH > 7: basic
Example:

  • A solution with a pH of 3 is acidic, while a solution with a pH of 9 is basic.

14.3 Calculating pH
Strong Acids: Fully dissociate in solution (i.e., HCl).
Weak Acids: Partially dissociate; ionization results in equilibrium.
Acid-Dissociation Constant: K_a = \frac{[H^+][A^-]}{[HA]}
Example:

  • For a 0.01 M HCl solution, since it is a strong acid, pH = 2. For acetic acid, the K_a allows for equilibrium concentration calculations.

14.4 Properties of Bases
Strong Bases: Fully dissociating bases, such as NaOH.
Weak Bases: Partially dissociating bases, like NH_3.
Example:

  • Sodium hydroxide (NaOH) completely dissociates in water, whereas ammonia (NH₃) only partially dissociates, resulting in a less basic solution.

14.5 Acid-Base Properties of Salts
Neutral Solutions: Salts from strong acids and bases (e.g., NaCl).
Basic Solutions: When anions are conjugate bases of weak acids (e.g., NaF).
Acidic Solutions: When cations are conjugate acids of weak bases (e.g., NH_4Cl).
Example:

  • The reaction of $NaCl$ in water results in a neutral solution, while $NH4Cl$ creates an acidic solution due to the presence of $NH4^+$ ions.

14.6 Additional Acid-Base Properties
Acidity increases with higher electronegativity and lower bond strength in nonmetal hydrides.
Example:

  • Hydrofluoric acid (HF) is a stronger acid than acetic acid (CH₃COOH) due to the electronegativity of fluorine.

14.7 Lewis Acid-Base Model
Acids accept electrons; bases donate electrons.
Example:

  • The reaction of boron trifluoride (BF₃) with ammonia (NH₃) demonstrates Lewis acid-base behavior; BF₃ is the Lewis acid as it accepts an electron pair from NH₃, the Lewis base.

14.8 Strategies for Solving Problems
Identify major species and reactions.
Use appropriate equations for equilibria.
Determine concentration changes and solve for [H^+] and pH.
Example:

  • In a titration problem, to find the final pH, identify the weak acid and strong base involved, then apply the corresponding equilibrium equations to compute the resulting pH after neutralization.