Acids and Bases Notes

Acids and Bases - Topic 4

Arrhenius Theory of Acids and Bases

• Proposed by Svante Arrhenius (1859-1927) to describe acid and base behavior in water.
• Acids and bases dissociate/ionize in aqueous solutions.

Acids (Arrhenius)

• A substance that ionizes in water to produce one or more hydrogen ions, H^+ (aq)
• Example: HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)

Bases (Arrhenius)

• A substance that dissociates in water to form one or more hydroxide ions, OH^- (aq)
• Example: NaOH(aq) \rightarrow Na^+(aq) + OH^-(aq)
• H^+ (aq) ion is responsible for the properties of acids.

Strengths of the Arrhenius Theory

• Explains why acids and bases are electrolytes.
• Explains why acids and bases have different chemical properties.

Limitations of the Arrhenius Theory

• Doesn't explain why some compounds are acidic or basic in water (e.g., NH3 is basic, CO2 is acidic).

Ionization vs. Dissociation

Ionization

• A reaction in which electrically neutral molecules or atoms produce ions:
• HCl(aq) + H2O(aq) \rightarrow H3O^+(aq) + Cl^-(aq)
• H_3O^+(aq) – hydronium ion, often written as H^+(aq) for simplicity.

Dissociation

• The process in which ions break apart when dissolved in solution.

Bronsted-Lowry Acids & Bases

• Developed in 1923 as a better definition for acids and bases.

Bronsted-Lowry Acid

• A substance that donates a hydrogen ion to any other substance; a proton donor.

Bronsted-Lowry Base

• A substance that accepts a hydrogen ion; a proton acceptor.
• Most accurate definition.

Conjugate Acids & Bases

• Chemical equations often reach equilibrium, where forward and reverse reactions occur at equal rates.
• The double arrow indicates equilibrium.
• Example: HCl(aq) + H2O(l) \rightleftharpoons H3O^+(aq) + Cl^-(aq)

Conjugate Acid-Base Pair

• Two molecules or ions related by the transfer of a proton.

Conjugate Base of an Acid

• The particle remaining when a proton is removed from the acid.
• Example: HCl (aq) \text{ (acid)} \xrightarrow{\text{remove proton}} Cl^- (aq) \text{ (conjugate base)}

Conjugate Acid of a Base

• The particle produced when a base receives a proton.
• Example: H2O (l) \text{ (base)} \xrightarrow{\text{add proton}} H3O^+ (aq) \text{ (conjugate acid)}
• In the reaction: HCl(aq) + H2O(l) \rightleftharpoons H3O^+(aq) + Cl^-(aq)
  • HCl(aq) (Acid) and Cl^-(aq) (Conjugate Base) form a conjugate acid-base pair.
  • H2O(l) (Base) and H3O^+(aq) (Conjugate Acid) form a conjugate acid-base pair.

Properties of Acids and Bases

The pH Scale

• pH stands for "potential hydrogen."
• Measures the amount of H^+ ions in a solution.
• Logarithmic scale: each pH unit change represents a factor of 10 change in concentration.
• Ranges from 0-14.
• Acidic solutions: pH < 7 (as acidity increases, pH decreases).
• Basic solutions: pH > 7 (as basicity increases, pH increases).

Examples for pH Scale

• How much more acidic is:
  • a pH of 2 than a pH of 3? (Answer: 10 times)
  • a pH of 2 than a pH of 5? (Answer: 1000 times)

Labelling the pH Scale

• Strongly acidic solutions: pH < 2.
• Weakly acidic solutions: 2 < pH < 7.
• Weakly basic solutions: 7 < pH < 12.
• Strongly basic solutions: pH > 12.
• Neutral solutions: pH = 7.

pH or pOH of a Solution

• Relationships:
  • [H^+(aq)] = 10^{-pH}
  • pH = -log[H^+(aq)]
  • [OH^-(aq)] = 10^{-pOH}
  • pOH = -log[OH^-(aq)]
• [ ] brackets mean molar concentration (mol/L).
• pH + pOH = 14

Calculations Involving pH

• Calculate the pH of a solution with:
  • [H^+] = 4.2 \times 10^{-4} \text{ mol/L}
  • [H^+] = 3.5 \times 10^{-8} \text{ mol/L}
• Calculate the [H^+] if the solution has a:
  • pH of 9.6
  • pH of 3.8
• Calculate the pOH if:
  • [H^+] = 3.1 \times 10^{-7} \text{ mol/L}

Calculations Involving pH - Answers

• Calculate the pH of a solution with:
  • [H^+] = 4.2 \times 10^{-4} \text{ mol/L}. Answer: 3.37
  • [H^+] = 3.5 \times 10^{-8} \text{ mol/L}. Answer: 7.46
• Calculate the [H^+] if the solution has a:
  • pH of 9.6. Answer: 2.51 \times 10^{-10} \text{mol/L}
  • pH of 3.8. Answer:1.58 \times 10^{-4} \text{mol/L}
• Calculate the pOH if:
  • [H^+] = 3.1 \times 10^{-7} \text{ mol/L}. Answer: 7.5

Acids with Multiple Protons

• Monoprotic acids: contain only one hydrogen (e.g., HNO_3).
• Diprotic acids: contain two hydrogens (e.g., H2SO4).
• Triprotic acids: contain three hydrogens (e.g., H3PO4).

Factors Affecting Acid and Base Strength

• Acid strength is determined by the concentration of H^+ ions in solution (higher concentration = stronger acid).
• Base strength is determined by the concentration of OH^- ions in solution (higher concentration = stronger base).
• The most important factor is the degree of ionization or dissociation.

Strong Acids

• Ionize 100% (>99%) upon dissolving in water.
• Very strong electrolytes due to the high concentration of ions.
• There are 6 common strong acids.

Weak Acids

• Transfer only a fraction of their protons to water; most remain intact.
• Partly ionized.
• Poor electrolytes due to the low concentration of ions.
• Includes most other acids that aren't strong acids.

Strong Bases

• Dissociate 100% upon dissolving in water.
• Very good electrolytes due to the high concentration of ions.

Weak Bases

• Only a fraction of molecules accept protons; partly ionized.
• Poor electrolytes due to the low concentration of ions.
• Example: ammonia.
• Includes other bases that aren't strong bases.

Concentration vs. Strength

• Strong ≠ concentrated.
• Weak ≠ diluted.
• Concentration refers to the number of moles in a volume.
• Strength refers to the degree of ionization.

Homework

• Read Section 10.1 - pg 454-463
• Practice Worksheet