Acids and Bases Chapter 14

Concepts of Acids and Bases

• Three primary definitions: Arrhenius, Brønsted-Lowry, and Lewis.
• Arrhenius Definition: Acids produce H3O+ in water; bases produce OH− in water.
• Brønsted-Lowry Definition: Acids donate H+, while bases accept H+ during reactions.
• Lewis Definition: Acids accept electron pairs; bases donate electron pairs.

Arrhenius Concept

• Strong acids like HCl fully dissociate in water: HCl → H+ + Cl−.
• Strong bases like NaOH fully dissociate in water: NaOH → Na+ + OH−.
• Acid-base reactions produce water and a salt: acid + base → salt + water (exothermic).

Brønsted-Lowry Theory

• Key focus on H+ transfer.
• Acid: proton donor; Base: proton acceptor.
• Reactions involving proton transfer define acid-base interactions.

Conjugate Acid-Base Pairs

• Pair consists of acid and base differing by one proton.
• Example: HCl (acid) and Cl− (conjugate base).

Brønsted-Lowry Acids and Bases

• Any substance with H can act as a Brønsted-Lowry acid.
• Brønsted-Lowry bases are substances with lone pairs to accept protons.

Lewis Acids and Bases

• Lewis acids are electron pair acceptors (electron deficient).
• Lewis bases are electron pair donors (must possess lone pairs).

Strong vs Weak Acids and Bases

• Strong acids/bases fully ionize in solution, acting as strong electrolytes.
• Weak acids/bases partially ionize, resulting in weak electrolytes.

Relative Strengths of Acids and Bases

• Strongest acids have weakest conjugate bases and vice versa.
• Equilibrium favors formation of weaker acid/base.

Molecular Structure Influencing Acid Strength

• Binary acids (H—Y): Acidity increases across a period due to electronegativity (e.g., HCl > HF).
• Acidity increases down a group as H—Y bond strength decreases.

Strengths of Oxoacids

• Acid strength increases with electronegativity of Y in H-O-Y- compounds.
• For oxoacids with varying O atoms, strength increases with the number of O atoms.

Autoionization of Water

• Pure water undergoes autoionization: 2 H2O ⇌ H3O+ + OH−.
• Concentration of H3O+ and OH− in pure water is 1 × 10−7 M at 25 °C.

Ion-Product Constant for Water (Kw)

• Kw = [H3O+][OH−] = 1.0 × 10−14 at 25 °C.
• Adding acid/base alters [H3O+] and [OH−], maintaining Kw constant but in opposite proportions.

pH Measurement

• pH is calculated as: pH = -log[H3O+].
• pH < 7 indicates acidic conditions; pH > 7 indicates basic conditions; pH = 7 indicates neutrality.

Solutions of a Strong Acid

• Strong acids contribute significantly to H3O+ concentration in solution.
• Calculating [OH−] involves considering autoionization: Kw = [H+][OH−].

Solutions of a Weak Acid

• Weak acids ionize partially, producing less H3O+.
• Their conjugate bases can contribute to the basicity of a solution.

Cations and Weak Acids

• Some cations derived from weak bases can act as acids.
• Cations from strong bases typically do not alter pH.

Assessing Salt Solution pH

• Neutral: salt from strong acid and strong base.
• Basic: salt from strong base and weak acid.
• Acidic: salt from weak base and strong acid.

Common Ion Effect

• Addition of a common ion shifts equilibrium to favor reactants, reducing ionization of weak acids.
• Example: Adding HCl to acetic acid reduces H3O+ production.

Buffer Solutions

• Buffers resist changes in pH when acid or base is added.
• Composed of weak acid and conjugate base or weak base and conjugate acid.
• Example: Blood buffer system involves H2CO3 and HCO3−.### Acidic Buffer Solution
• An acidic buffer requires significant amounts of a weak acid and its conjugate base for effective buffering.
• Key components: acetic acid (HAc) and sodium acetate (NaAc).
• Addition of strong acid (HCl) or strong base (NaOH) leads to neutralization reactions without significant pH changes if the quantities are appropriate.

Buffer Dynamics

• Strong Base Neutralization:
  • NaOH neutralizes weak acid (HC2H3O2) to form water and sodium acetate.
• Strong Acid Neutralization:
  • HCl neutralizes conjugate base (NaC2H3O2) to form acetic acid and sodium chloride.

Henderson-Hasselbalch Equation

• pH = pKa + log([A-]/[HA])
• Utilizes concentrations of the conjugate base ([A-]) and weak acid ([HA]) to calculate the pH of a buffer solution.
• Valid under the "x is small" approximation.

Buffer Effectiveness and Capacity

• Effective buffers can neutralize moderate amounts of acid or base without significant pH changes.
• Buffering range: Effective when 0.1 < [base]:[acid] < 10.
• Buffering capacity increases with the concentration of acid and base components.

Choosing a Buffer

• Select an acid with a pKa close to the desired pH for optimal buffering.

Titration Process

• A titration involves adding a titrant of unknown concentration from a burette to a known solution until the reaction reaches the endpoint.
• Endpoints may be indicated by pH-sensitive color changes (indicators).
• Equivalence point occurs when moles of H3O+ equal moles of OH−.

Titration Curves

• Plotting pH against titrant volume shows shifts, with the inflection point marking the equivalence point.
• pH at equivalence depends on the nature of the salt produced:
  • Neutral salt: pH = 7
  • Acidic salt: pH < 7
  • Basic salt: pH > 7

Buffering with Different Acids and Bases

• Buffer regions are critical for weak acids versus strong bases and vice versa, demonstrating varying pH responses to added titrant.
• Equal concentrations of acid and base enhance buffer effectiveness.

Summary of Buffer Behavior

• Buffering capacity contributes significantly to pH stability.
• Effective when both components are present in similar concentrations, optimizing the ability to neutralize added acids or bases.

Acid-Base Chemistry Overview

• The concentration of hydronium ions [H3O+] and hydroxide ions [OH–] in pure water is equal to 1.00 × 10−7 M at 25 °C.
• Acidic solutions: [H3O+] > 1.00 × 10−7 M, [OH–] < 1.00 × 10−7 M.
• Basic solutions: [OH–] > 1.00 × 10−7 M, [H3O+] < 1.00 × 10−7 M.

pH Concept

• pH indicates the acidity or basicity of a solution and is calculated using the formula: pH = -log[H3O+].
• A solution with [H3O+] of 1.0 x 10-3 M has a pH of 3.
• pH scale:
  • pH < 7 indicates acidity
  • pH > 7 indicates basicity
  • pH = 7 indicates neutrality
• Normal pH range is 0 to 14.

Hydroxide Ion Concentration (pOH)

• pOH can be calculated using the formula: pOH = -log[OH–].
• The relationship between pH and pOH is given by: pH + pOH = 14.
• At 25 °C, the ion product of water (Kw) is: Kw = [H3O+][OH–] = 1.0 × 10−14.

Strong Acids and Bases

• Strong acids dissociate almost completely in water, contributing significantly to H3O+ concentration.
• Strong bases contribute to OH– concentration similarly.
• In a solution of a strong acid, the contribution from water is negligible as [H3O+] significantly outweighs [OH–].
• The ion product remains constant, affecting inverse relationships between [H3O+] and [OH–].

Examples of Strong Acids and Bases

• For 0.10 M HCl, the hydroxide concentration can be determined using the ion product constant:
  • Kw = [H+] [OH–].
• Calculating H3O+ and OH– concentrations for specific examples:
  • 0.15 M HNO3 leads to specific H3O+ and OH– values.
  • 0.010 M Ca(OH)2 can also be used for calculations.

Characteristics of Acid Strength

• Acid strength increases with the increase in atom size and decreases in H—X bond strength down a group.
• Binary acid strength: H—C < H—N < H—O < H—F.
• Acid strength trends for oxoacids:
  • Strength increases with the electronegativity of atom Y in H—O—Y.
  • Example: HIO < HBrO < HClO.
  • Strength increases with the number of oxygen atoms bonded to Y: HClO < HClO2 < HClO3 < HClO4.

Autoionization of Water

• Pure water behaves as a nonelectrolyte and minimally conducts electricity.
• Autoionization reaction:
  • H2O(l) + H2O(l) ⇌ H3O+(aq) + OH−(aq).
• In aqueous solutions, concentrations of H3O+ and OH– are equal: [H3O+] = [OH–] = 10−7 M at 25 °C.

Equilibrium Concepts

• The equilibrium constant expression for the autoionization of water can be represented as:
  • Kw = [H3O+][OH–].
• The equilibrium constant Kw is consistently 1.0 × 10−14 at 25 °C, indicating the relationship between H3O+ and OH– concentrations in water.

Learn about strong acids and bases, Arrhenius, Bronsted-Lowry, and Lewis definitions. Identify common acids and bases with examples.