Chem 2 Chapter 16 Questions (medium)

Introduction to Acids, Bases, and pH

• Arrhenius acids yield H⁺ ions in water; Arrhenius bases yield OH⁻ ions in water.
• Brønsted-Lowry acids are proton donors; Brønsted-Lowry bases are proton acceptors.
• A conjugate acid-base pair contains two substances differing by one proton (H⁺).
• HF and F⁻ are a conjugate acid-base pair; HF donates a proton becoming F⁻, and F⁻ accepts a proton becoming HF.
• The conjugate base for NH₃ is NH₂⁻.
• The conjugate acid for NH₃ is NH₄⁺.
• Writing a conjugate base means the original species acts as an acid, donating a proton.
• Writing a conjugate acid means the original species acts as a base, accepting a proton.
• Autoionization of water reaction: H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)
• Kw expression: Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C
• Kw is an equilibrium constant expression based on products over reactants for water's autoionization.
• Amphoteric substances can act as both acids and bases.
• Water (H₂O) is amphoteric, donating or accepting a proton.
• H⁺ is a proton; in water, it combines with H₂O to form H₃O⁺ (hydronium ion).
• H⁺ and H₃O⁺ are often used interchangeably to refer to hydrogen ions in solution.
• Kw is the product of [H₃O⁺] and [OH⁻] in water and the pH scale is based on [H₃O⁺].
• Increased [H₃O⁺] decreases pH; decreased [H₃O⁺] increases pH.
• A neutral solution has a pH of 7, where [H₃O⁺] = [OH⁻].
• The pH scale is technically accurate at 25°C.
• Acidic solutions have a pH < 7; basic solutions have a pH > 7.

Strong Acids/Bases vs Weak Acids/Bases

• Strong acids fully dissociate, producing H⁺ ions, allowing pH calculation via the negative logarithm of the strong acid's concentration ([HCl] = [H⁺] = [H₃O⁺]).
• With strong bases like Ba(OH)₂, [OH⁻] is twice the concentration of Ba(OH)₂: Ba(OH)₂ → Ba²⁺ + 2OH⁻, and pH is found using pOH = -log[OH⁻], then pH = 14 - pOH.
• Seven strong acids: HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄, HClO₃.
• Strong bases completely dissociate in water, releasing OH⁻ ions and include Group 1 metal hydroxides like NaOH and KOH, and some Group 2 metal hydroxides like Ba(OH)₂.
• 'p' function meaning: the negative logarithm of; pH = -log[H⁺], pOH = -log[OH⁻].
• Weak acids or bases only partially dissociate in water, lowering their effectiveness in proton donation/acceptance, unlike the 100% dissociation of strong acids/bases.
• For weak acids: HA + H₂O ⇌ H₃O⁺ + A⁻, with equilibrium expression Ka = [H₃O⁺][A⁻] / [HA].
• For weak bases: B + H₂O ⇌ BH⁺ + OH⁻, with equilibrium expression Kb = [BH⁺][OH⁻] / [B].
• Weak acid/base pH calculation involves using the Ka/Kb expression and an ICE table to determine ion concentrations at equilibrium.
• Strong acids/bases fully dissociate, allowing pH calculation from initial concentrations without needing an ICE table.
• ICE chart chemical equations: acid + water ⇌ hydronium ion + conjugate base, or base + water ⇌ hydroxide ion + conjugate acid.
• ICE chart problems involve finding the value of the equilibrium constant, or using the value to find pH and concentration.
• pH equations link concentrations and pH/pOH values, and the approximation method can simplify solving quadratic equations.
• An acid's strength is inversely related to its conjugate base's strength: stronger acids form weaker conjugate bases, and stronger bases form weaker conjugate acids.
• Strong acid/base conjugates are very weak and have negligible acid/base properties.
• For a conjugate acid-base pair, Kw = Ka × Kb; high Ka means low Kb, and vice versa.

Acids, Bases, and Salt Solutions

• Diprotic acids like H₂SO₄ can donate two protons (H⁺).
• Polyprotic acids, like H₃PO₄, can donate more than one proton.
• HF is monoprotic.
• Salt solution pH depends on the strength of the originating acid and base.
• A salt from a strong acid and weak base is acidic; a salt from a weak acid and strong base is basic; a salt from a strong acid and strong base is neutral.
• Conjugates of strong acids/bases (like Cl⁻) are weak and considered pH-neutral.
• For a solution with strong and weak acids, ignore the hydronium ions from the weak acid because the strong acid dominates the pH.
• For a solution with strong and weak bases, ignore the hydroxide ions from the weak base because the strong base dominates the pH.
• NH₄Cl hydrolysis: NH₄Cl (s) → NH₄⁺ (aq) + Cl⁻ (aq).
• NH₄⁺ acts as a weak acid, reacting with water: NH₄⁺ (aq) + H₂O (l) ⇌ NH₃ (aq) + H₃O⁺ (aq), which acidifies the solution.
• Cl⁻ doesn't affect pH because it's from a strong acid (HCl) and doesn't react with water.
• Lewis acids accept electron pairs, and Lewis bases donate them.
• Arrhenius definition is most specific, Lewis definition is broadest; all Arrhenius acids are Lewis acids, but not all Lewis acids are Arrhenius acids.