Notes on Acids and Bases

Chapter 19: Acids & Bases

General Properties of Acids

• Taste: Sour
• Feel: Wet
• pH: Less than 7 (range: 0-6)
• Litmus Test: Turns litmus paper red
• Phenolphthalein: Colorless in acid solutions
• Reactions:
  • React with metals to produce hydrogen gas ($H_2$)
  • React with carbonates to produce carbon dioxide gas ($CO_2$)
• Conductivity: Conduct electricity

General Properties of Bases

• Taste: Bitter
• Feel: Slippery
• pH: Greater than 7 (range: 8-14)
• Litmus Test: Turns litmus paper blue
• Phenolphthalein: Pink in base solutions
• Reactions:
  • Do not react to form hydrogen gas ($H_2$)
• Conductivity: Conduct electricity

Naming Acids

• For Halogen Anions: Use the prefix "hydro" and change the ending to "ic".
  • Example: $HCl$ → Hydrochloric Acid
• For Polyatomic Anions: Name changes based on the ending:
  • Ends in ATE → becomes IC
  • Example: $HNO_3$ → Nitric Acid
  • Ends in ITE → becomes OUS
  • Example: $HNO_2$ → Nitrous Acid

Common Acids

• Hydroiodic acid ($HI$)
• Hydrobromic acid ($HBr$)
• Hydrofluoric acid ($HF$)
• Sulfuric acid ($H<em>2SO</em>4$)
• Carbonic acid ($H<em>2CO</em>3$)
• Acetic acid ($H(C<em>2H</em>3O_2)$)
• Sulfurous acid ($H<em>2SO</em>3$)

Naming Bases

• Bases use standard ionic naming rules.

Common Bases

• Sodium hydroxide ($NaOH$)
• Potassium hydroxide ($KOH$)
• Lithium hydroxide ($LiOH$)
• Calcium hydroxide ($Ca(OH)_2$)
• Barium hydroxide ($Ba(OH)_2$)
• Magnesium hydroxide ($Mg(OH)_2$)

Arrhenius Model

• Acids: Substances that contain hydrogen and ionize to produce $H^+$ ions.
  • Example: $HCl(g) → H^+(aq) + Cl^-(aq)$
• Bases: Substances that contain hydroxide and dissociate to produce $OH^-$ ions.
  • Example: $NaOH(s) → Na^+(aq) + OH^-(aq)$
• Neutral Solutions: Contain equal numbers of $H^+$ and $OH^-$ ions ($H_2O ↔ H^+ + OH^-$)

Bronsted-Lowry Model

• Acid: Substance that donates a hydrogen ion ($H^+$).
• Base: Substance that accepts a hydrogen ion ($H^+$).

Conjugate Acid-Base Pair

• Comprises two substances related by the donation and acceptance of a single hydrogen ion.
• An acid has one extra $H^+$ than its conjugate base.

Amphoteric Substances

• Substances like water that can act as both acids and bases.

Monoprotic & Polyprotic Acids

• Monoprotic Acids: Have one hydrogen to donate (e.g., $HCl$, $HNO_3$).
• Polyprotic Acids: Have more than one hydrogen to donate (e.g., $H<em>2SO</em>4$, $H<em>3PO</em>4$).

Ionization of Polyprotic Acids

• Each hydrogen ion lost occurs in separate reactions, e.g.,
  • $H<em>2PO</em>4(aq) + H<em>2O(l) ⇌ H</em>3O^+(aq) + HPO_4^{2-}(aq)$
  • $HPO<em>4^{2-}(aq) + H</em>2O(l) ⇌ H<em>3O^+(aq) + PO</em>4^{3-}(aq)$

Strength of Acids & Bases

• Strong Acids: Ionize completely (e.g., $HCl$, $H<em>2SO</em>4$, $HNO_3$).
• Weak Acids: Partially ionize (e.g., acetic acid $HC<em>2H</em>3O_2$).
• Strong Bases: Dissociate entirely into metal ions and hydroxide ions (e.g., $NaOH$, $KOH$, $Ca(OH)_2$).
• Weak Base: Example includes ammonia ($NH_3$).

The pH Scale

• Ranges from 0 to 14.
• Neutral solution: pH = 7 ($[H^+] = [OH^-]$).
• Acidic solution: pH < 7 ($[H^+] > [OH^-]$).
• Basic solution: pH > 7 ($[H^+] < [OH^-]$).

pH Calculations

• Formulas:
  • $pH = - ext{log} [H^+]$
  • $pOH = - ext{log} [OH^-]$
  • $pH + pOH = 14$
  • $[H^+][OH^-] = 1 imes 10^{-14}$

Neutralization

• Reaction: Acid and base react in aqueous solution to produce a salt and water.
  • Example: $HCl + NaOH → NaCl + H_2O$
• Indicators in Titration: Indicators change color at the equivalence point but not always at pH 7.

Buffered Solutions

• Buffers: Solutions that resist changes in pH when limited amounts of acid or base are added, typically a mixture of a weak acid and its conjugate base or vice versa.