Acid-Base Chemistry: Definitions and Properties

Acid-Base Basics

• Arrhenius acids produce H⁺ in water, while bases produce OH⁻.
• Brønsted-Lowry acids are H⁺ donors, and bases are H⁺ acceptors.
• Lewis acids accept electron pairs, and Lewis bases donate electron pairs.
• Binary acids contain two elements: a cation (H⁺) and a monatomic anion, named with a hydro- prefix and an -ic suffix.
• Oxyacids contain hydrogen, oxygen, and another element; the naming is based on the polyatomic anion: -ate becomes -ic, and -ite becomes -ous.
• Acids taste sour, dissolve some metals, turn litmus paper red, and neutralize bases.
• Bases feel slippery, taste bitter, turn litmus paper blue, and neutralize acids.
• Strong acids/bases dissociate completely in water, while weak acids/bases dissociate partially.

Acid-Base Ionization and K

• Ka is the acid ionization constant for the reaction HA(aq) + H₂O(l) ⇋ H₃O⁺(aq) + A⁻(aq), given by Ka = [H₃O⁺][A⁻]/[HA].
• A larger Ka means a stronger acid.
• Kb is the base ionization constant for the reaction B(aq) + H₂O(l) ⇋ BH⁺(aq) + OH⁻(aq), given by Kb = [BH⁺][OH⁻]/[B],
• A larger Kb indicates a stronger base.
• Kw is the ion product for water, Kw = [H₃O⁺][OH⁻] = 1 x 10⁻¹⁴ at 25 °C.
• In neutral solutions, [H₃O⁺] = [OH⁻] = 1 x 10⁻⁷ at 25°C, with pH = 7.
• Acidic conditions are when [H₃O⁺] > [OH⁻] (pH < 7), and basic conditions are when [H₃O⁺] < [OH⁻] (pH > 7).

pH and pOH

• pH = -log[H⁺], the number of decimal places should match the significant figures in the original concentration.
• pOH = -log[OH⁻].
• pH + pOH = 14 at 25°C.

Solving Acid-Base Problems

• For strong acids/bases that dissociate completely, all of the initial acid/base is converted to H⁺/OH⁻, can be used to find pH.
• For weak acids/bases that dissociate partially, use an ICE table and the Ka/Kb expression to determine species concentrations at equilibrium.
• 5% rule: if Ka is small and the initial concentration of acid is not small, x can be approximated as negligible in the denominator.

Salts and Ions

• Anions of strong acids are neutral, while anions of weak acids are basic.
• Cations of strong bases are neutral, while cations of weak bases are acidic.
• Compare the relative strengths of weak acids by examining their Ka values; the acid with the higher Ka will produce a stronger conjugate base.

Polyprotic Acids

• Polyprotic acids have multiple ionizable protons and dissociate in successive steps.
• Each dissociation step has its own Ka value that decreases with each step.
• When calculating the pH of a polyprotic acid solution, generally can be simplified by only considering the first dissociation step.

Acid Strength and Molecular Structure

• For binary acids (H-X), if X is highly electronegative, the H⁺ cannot be easily removed and is weak. If X is not very electronegative, H⁺ CAN be removed, acid is STRONG.
• For oxyacids (H-O-Y), the more electronegative Y is, the easier it is to remove H⁺, and thus the acid is stronger.

Buffers

• Buffers resist changes in pH upon addition of small amounts of acid or base.
• They contain a weak acid and its salt or a weak base and its salt.
• Buffers are most effective when the concentrations of both components are large and equal.
• The buffer range is the pH range over which the buffer is effective, with the pKa of the buffer close to the desired pH.
• The concentrations of the buffer components should not differ by more than a factor of 10.
• Henderson-Hasselbalch equation is used to calculate the pH of a buffer solution: pH = pKa + log([A⁻]/[HA])

Titration and pH Curves

• Titration involves plotting pH vs. volume of titrant added.
• Equivalence point: the point where moles of H⁺ equals moles of OH⁻.
• This is where only water and the salt produced.
• Indicator: a substance that changes color near the equivalence point, with a pKa +/- 1 from the pH at the equivalence point
• Endpoint: the point where the indicator changes color, close to the equivalence point.
• Strong Acid/Strong Base: calculate 5 points
  • At beginning - no titrant added
  • At equivalence point pH = 7.
  • 1 mL before the equivalence point
  • 1 mL after the equivalence point
  • Approximately 10 mL after equivalence point
• Weak Acid/Strong Base or Weak Base/Strong Acid - calculate 6 points
  • At beginning - no titrant added
  • At equivalence point pH will not be 7.
  • Half-equivalence point. Here the pH=pKa - that's your y value.
  • 1 mL before the equivalence point
  • ImL after the equivalence point
  • Approximately 10 mL after equivalence point

Acid-Base Indicators

• Serve to identify the Equivalence Point.
• Consist of a weak acid and conjugate base or vice versa that are organic and complex with different colors for the acid and conjugate.

Solubility Equilibria

• Ksp (solubility product constant) describes the equilibrium between a solid and its ions in a saturated solution.
• Write the equation and expression as follows: (solid<->ions)
• Equation of Ksp = [ion][ion]
• The process of dissociation is:
  • Solute is added to water.
  • Solute (ionic) dissociates to some small degree
  • Those ions can then precipitate.
  • solution which is saturated!!!
• Experimentally determined solubility (in M or any unit) can be used to find Ksp.
• If Ksp is known, solubility (any unit) can be determined
• Compare Solublity with Ksp, use Ksp (greater Ksp, the more soluble, as a greater concentration of ions will be present in the saturated solution)
• Common Ion Effect basically states Le Chatelier's Principle! Where a species will reduce the solubility of your solute by stress in the reaction.
• pH < 7 or if any acid is added, will that acid react with the anion of the equation?
• Selective Precipitation:
  • A species can be chosen that will only precipitate the desired ion and not the others.
  • A species can be chosen that will precipitate one ion before the others.
    • Let's call this [Species]min
    • Let's call this [Species]max