Acids and Bases

Acids and Bases - Definitions

  • Arrhenius Definition
    • Acid: A species that dissociates in water to produce hydrogen ions ($H^+$ or H3O^+$).
    • Example:
      • HCl(aq)+H2O(l)→Cl−(aq)+H3O+(aq)HCl(aq) + H2O(l) → Cl^-(aq) + H3O^+(aq)
    • Base: A species that dissociates in water to produce hydroxide ions ($OH^-$).
    • Example:
      • NaOH(s)+H2O(l)→Na+(aq)+OH−(aq)NaOH(s) + H2O(l) → Na^+(aq) + OH^-(aq)
  • Founder of Theory: Svante August Arrhenius, Nobel Prize in Chemistry 1903.

Acid-Base Reactions

  • General Reaction Type:
    • Involve transfer of a hydrogen ion ($H^+$) from an acid to a base.
    • Example:
    • HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
      • Here, H+H^+ from HClHCl is transferred to OH−OH^- from NaOHNaOH, forming water.
  • Product Formation: Always produces water, confirming Arrhenius' theory.

Proton Donors and Acceptors

  • Brønsted-Lowry Theory
    • Acid: Proton donor.
    • Example:
      • HCl(aq)+H2O(l)→Cl−(aq)+H3O+(aq)HCl(aq) + H2O(l) → Cl^-(aq) + H3O^+(aq)
    • Base: Proton acceptor.
    • Example:
      • NH3+H2O→NH4++OH−(aq)NH3 + H2O → NH4^+ + OH^-(aq)

Conjugate Acid-Base Pairs

  • Definitions:
    • Conjugate Acid: The species formed when a base accepts a proton.
    • Conjugate Base: The species formed when an acid donates a proton.
    • Conjugate Pair: Acid and base pairs related by this transfer.

Strength of Acids and Bases

  • Strong Acids/Bases:
    • Complete Dissociation in water.
    • Example:
      • HCl(g)+H2O(l)→H+(aq)+Cl−(aq)HCl(g) + H2O(l) → H^+(aq) + Cl^-(aq)
    • Strong acids: HCl, HNO3, H2SO4, etc.
    • Weak Acids/Bases:
    • Partially dissociate in solution, achieving equilibrium.
    • Example:
      • CH3COOH(aq)+H2O(l)⇌CH3COO−(aq)+H3O+(aq)CH3COOH(aq) + H2O(l) ⇌ CH3COO^-(aq) + H3O^+(aq)

Dissociation of Water

  • Autoionization:
    • Water can act both as acid and base, producing H3O+H3O^+ and OH−OH^- by dissociation.
    • H2O(l)+H2O(l)⇌H3O+(aq)+OH−(aq)H2O(l) + H2O(l) ⇌ H3O^+(aq) + OH^-(aq)
  • Ion Product Constant ():
    • Kw=[H3O+][OH−]=1.0imes10−14K_w = [H3O^+][OH^-] = 1.0 imes 10^{-14} at 25°C.

Measuring pH

  • pH Scale Range:
    • 0 (acidic) to 14 (basic).
  • pH Calculation:
    • Given [H3O+][H3O^+]
    • pH=−extlog([H3O+])pH = - ext{log}([H3O^+])
  • pOH and their Relation:
    • pOH=−extlog([OH−])pOH = - ext{log}([OH^-])
    • pH+pOH=14pH + pOH = 14

Titration and Buffer Solutions

  • Buffer Solutions:
    • Combinations of weak acids/bases and their conjugate bases/acids, resisting changes in pH when small amounts of acids/bases are added.
  • Henderson-Hasselbalch Equation:
    • pH=pKa+extlog[A−][HA]pH = pKa + ext{log}\frac{[A^-]}{[HA]}
    • Useful for calculating pH of buffer systems.
  • Titration Curves:
    • Reveal reaction progress in acid-base neutralization, showing equivalence points and buffer regions.
    • Types of Curves:
    • Strong acid with strong base: sharp rise at equivalence point.
    • Weak acid with strong base: inflection and buffer regions present.

Example Problems for pH Calculation

  • Given Concentrations:
    • From strengths and calculated equilibrium conditions, find pH.
  • Polyprotic Acids:
    • May involve successive deprotonation, requiring separate equilibrium considerations for each proton.
    • Example problem: H2SO4(aq) + 2NaOH(aq)
      ightarrow Na2SO4(aq) + 2H_2O(l)

Acidity of Hydrated Metal Ions

  • Acidity behavior:
    • Hydrated metal cations (Fe$^{3+}$, Al$^{3+}$) can interact with water to donate protons, contributing to an acidic environment.

Determining Acid-Base Behavior in Ion Compounds

  • Assessing compounds:
    • Classification as acidic or basic.
    • Consider cations from strong acids and anions from weak acids, their behavior in dissolving.
    • Examples:
    • Strong acid + strong base → neutral solution,
    • Weak acid + strong base → basic solution due to remaining OH-

Conclusion

  • Acid-Base Theories:
  • Arrhenius and Brønsted-Lowry definitions provide frameworks for understanding behavior in aqueous solutions, important in both theoretical chemistry and practical applications.