Acids and Bases Summary

Introduction to Acids and Bases

Arrhenius Model

• Acids produce hydrogen ($H^+$) ions in water.
• Bases produce hydroxide ($OH^−$) ions in water.
• Example:
  • $HCl(g) \rightarrow H^+(aq) + Cl^−(aq)$
  • $NaOH(s) \rightarrow Na^+(aq) + OH^−(aq)$

Brønsted–Lowry Definition

• Acid: Proton ($H^+$) donor.
• Base: Proton ($H^+$) acceptor.
• Example: $HCl(g) + H<em>2O(l) \rightarrow H</em>3O^+(aq) + Cl^−(aq)$
  • $HCl$ is a Brønsted–Lowry acid.
  • $H_2O$ is a Brønsted–Lowry base.

Brønsted–Lowry Acids

• Must contain a hydrogen atom.
• Common examples:
  • Hydrochloric acid ($HCl$)
  • Sulfuric acid ($H<em>2SO</em>4$)
  • Hydrobromic acid ($HBr$)
  • Nitric acid ($HNO_3$)

Naming Acids

• Anions ending in -ide: Add hydro- prefix and change -ide to -ic acid (e.g., $Cl^− \rightarrow HCl$: chloride to hydrochloric acid).
• Polyatomic anions ending in -ate: Change -ate to -ic acid (e.g., $SO<em>4^{2-} \rightarrow H</em>2SO_4$: sulfate to sulfuric acid).
• Polyatomic anions ending in -ite: Change -ite to -ous acid (e.g., $SO<em>3^{2-} \rightarrow H</em>2SO_3$: sulfite to sulfurous acid).

Brønsted–Lowry Bases

• Proton acceptors; must be able to form a bond to a proton.

Conjugate Acid-Base Pairs

• Differ by only one proton.
• Examples:
  • $HBr$ and $Br^−$
  • $H<em>2O$ and $H</em>3O^+$

Amphoteric Substances

• Can act as either an acid or a base.
• Water ($H<em>2O$) and bicarbonate ion ($HCO</em>3^−$) are common examples.
• $HCO<em>3^−(aq) + OH^−(aq) \rightarrow CO</em>3^{2−}(aq) + H_2O(l)$
• $HCO<em>3^−(aq) + H</em>3O^+(aq) \rightarrow H<em>2CO</em>3(aq) + H_2O(l)$

Acid and Base Strength

• Strong acids/bases: Ionize completely.
• Weak acids/bases: Ionize to a small extent.
• $HCl(g) + H<em>2O(l) \rightarrow H</em>3O^+(aq) + Cl^−(aq)$

Strong Acids

• Hydrochloric acid ($HCl$)
• Hydrobromic acid ($HBr$)
• Hydroiodic acid ($HI$)
• Nitric acid ($HNO_3$)
• Chloric acid ($HClO_3$)
• Perchloric acid ($HClO_4$)
• Sulfuric acid ($H<em>2SO</em>4$) (first ionization)

Weak Acids

• Acetic acid ($CH_3COOH$)
• Carbonic acid ($H<em>2CO</em>3$)
• Citric acid ($H<em>3C</em>6H<em>5O</em>7$)
• Hydrofluoric acid ($HF$)
• Hypochlorous acid ($HOCl$)
• Lactic acid ($HC<em>3H</em>5O_3$)
• Malic acid ($H<em>4C</em>4O_5$)
• Oxalic acid ($H<em>2C</em>2O_4$)
• Phosphoric acid ($H<em>3PO</em>4$)
• Tartaric acid ($H<em>2C</em>4H<em>4O</em>6$)

Polyprotic Acids

• Acids with more than one acidic hydrogen.
• Donate one $H^+$ ion at a time in steps.

Strong Bases

• LiOH, NaOH, KOH, RbOH, CsOH, $Ca(OH)<em>2$, $Sr(OH)</em>2$, $Ba(OH)_2$

Weak Bases

• Ammonia ($NH_3$)
• Methylamine ($CH<em>3NH</em>2$)
• Trimethylamine ($(CH<em>3)</em>3N$)

Ion-Product Constant of Water

• $K<em>w = [H</em>3O^+][OH^−]$

• At 25°C: $K<em>w = 1.0 × 10^{−14}$; in neutral solution, $[H</em>3O^+] = [OH^−] = 1.0 × 10^{−7} M$

• Neutral: $[H_3O^+] = [OH^−]$

• Acidic: [H_3O^+] > [OH^−]

• Basic: [OH^−] > [H_3O^+]

Calculating $[H_3O^+]$ and $[OH^−]$

• If $[H<em>3O^+]$ is known: $[OH^−] = \frac{1.0 × 10^{−14}}{[H</em>3O^+]}$
• If $[OH^−]$ is known: $[H_3O^+] = \frac{1.0 × 10^{−14}}{[OH^−]}$

The pH Scale

• $pH = −log[H_3O^+]$
• Acidic: pH < 7, $[H_3O^+] > 1 × 10^{−7} M$
• Neutral: pH = 7, $[H_3O^+] = 1 × 10^{−7} M$
• Basic: pH > 7, [H_3O^+] < 1 × 10^{−7} M

Calculating pH

• $pH = −log[H_3O^+]$

Calculating [H3O+] from pH

• $[H_3O^+] = 10^{−pH}$

Neutralization Reaction

• HA(aq) + MOH(aq) → $H_2O$ (l) + MA(aq) (acid + base -> water + salt)

Titration

• Used to determine unknown molarity.