Chapter 21 - Further Equilibria & Buffer Solutions

Kw

• Ionic Product of Water

• 1.00e-14 mol^2 dm^-3

• At 298K

Kc

• Equilibrium Constant

• Kc = [H+] [A-] / [HA]

• Since [HA] is constant for solids or salts and water, Kc = [H+] [A-]

Formula for pH of Strong Acids

• -log( [H+] )

• Assume all of the hydrogen ions come from the acid and not the water.

Formula for pH of Strong Bases

• -log( Kw/[OH-] )

• Assume complete dissociation, so the [OH-] is all of the OH- from the base.

• Kw = [H+] [OH-], so [H+] = Kw/[OH]

Ka

• Acid Dissociation Constant

• Higher = equilibrium to the right = complete dissociation

• Lower = equilibrium to the left = partial dissociation

• HA ←→ H + A

pKa

• -log(Ka)

Formula for Ka of Weak Acids

• Ka = [H+]² / [HA]

• Ignore the [H+] from the water as it is negligible.

• Assume the [HA] is the same as the concentration of acid because so little dissociates in solution.

Formula for pH for Weak Acids

• -log( sqrt(Ka * [HA]) )

• Utilize the Ka formula for Weak Acids to rearrange and solve for [H+]

Acid-Base Indicators

• Changes color over a specific pH range.

  • Bromothymol Blue (Yellow to Blue, pH 6.0-7.6)

  • Phenolphthalein (Colorless to Pink, pH 8.2-10.0)

  • Methyl Red (Red to Yellow, pH 4.2-6.3)

End Point

• The median of the pH range for acid-base indicators.

• or, the volume of alkali or acid added in a acid-base neutralization titration curve when neutralized

Acid-Base Neutralization Titration Curves

• In terms of Acid-Base:

  • Strong-Strong: Largest pH range

  • Strong-Weak: Normal pH range

  • Weak-Strong: Normal pH range

  • Weak-Weak: Smallest pH range (undetectable by indicators reliably)

Half-Neutralization Point

• Amount (moles) of alkali is equal to acid

• Exactly half of the volume from the End Point

• Kc = [H+] [A] / [HA], but [A] = [HA] so Kc = [H+]

• Helpful in calculating pH as Ka = [H+] or pKa = pH

Buffer Solution

• Solution where the pH is resistant to change when small amount of acid or alkali is added

• High amounts of conjugate acid and base to neutralize added ions

• Usually weak acid + its salt:

  • CH3OOH (Acid) ←→ CH3OO- (From NaCH3OO or Salt) + H+

Formula for pH of Buffer Solutions

• [H+] = Ka * [acid] / [salt]

• pH = pKa + log( [salt] / [acid] )

Buffer Solution in Blood

• HCO3-, Hemoglobin, Plasma, [H2PO4]-, [HPO4]2-

• HCO3- formed from CO2 dissolving in water, turns back into CO2 near the lungs.

• H+ increase = more water is created from HCO3-

• H+ decrease = Carbonic Anhydrase produces H2O + CO2 produces H+ and HCO3-

Solubility

• Amount of solute that dissolves in 1 dm³ of a saturated solution.

• mol dm³ (molar solubility) or g dm³ (mass solubility)

Saturated Solution

• Solution that contains as much dissolved solute as it can at a certain temperature.

Ksp

• Solubility Product

• Ksp = [ Cation ]^ # of moles * [ Anion] ^ # of moles

• Smaller Ksp = Less Soluble.

• Predicts PPT formation:

  • Calculated Ksp < Theoretical Ksp = No PPT

  • Calculated Ksp > Theoretical Ksp = PPT

Common Ion Effect

• Reduction in the solubility of a salt by adding a dissolved salt that shares an ion in both salts.

• Ag + Cl ←→ AgCl. Adding NaCl increase [Cl], and following le Chatelier’s Principle it shifts right.

Kpc

• Partition Constant

• Equilibrium Constant of the concentrations of a solute in 2 immiscible solvents that are next to each other at a certain temperature

• Kpc = [Solvent A] / [Solvent B]