Understanding Solutions: Saturated, Unsaturated

Types of Solutions

• A saturated solution contains the maximum amount of solute at a given temperature.
• An unsaturated solution contains less solute than it can dissolve.
• A supersaturated solution contains more solute than it can dissolve.

Factors Affecting Dissolving Rate

• Temperature: Increased entropy and enthalpy increase the dissolving rate.
• Surface Area: Higher surface area leads to faster dissolving.
• Stirring: Increased movement increases dissolving rate.
• Concentrated solutions have a lot of solute in a small amount of solvent.
• Dilute solutions have a small amount of solute in a large amount of solvent.
• Vinegar is a dilute solution, typically 95% H₂O and 5% Acetic Acid, requiring dilution to reach 100% concentration.
• More heat dissolves faster because entropy breaks down bonds.

Molarity

• The unit for concentration, defined as the number of moles of solute in 1 liter of solution.
• Molarity = Moles / Volume (in Liters).
• Moles = Molarity x Volume.
• Mass (g) = Volume (L) x Molarity (M) x Molar Mass (g/mol).
• Moles = mass / molar mass.

Dilution

• MsVs = MdVd, where:
  • Ms = Molarity of Stock Solution
  • Vs = Volume of Stock Solution
  • Md = Molarity of Dilute Solution
  • Vd = Volume of Dilute Solution
• Example: To prepare 500mL of 1.77M H₂SO₄ from an 18.0M stock solution, calculate the required volume of the stock solution. Vs = (MdVd) / Ms = (1.77M x 0.5L) / 18.0M = 0.00492L (49.2mL).
• Molarity is generally moles over mass.

Dynamic Equilibrium

• Irreversible Reaction: A + B → C + D
• Reversible Reaction: A + B ⇌ C + D
• Rate of forward reaction equals the rate of reverse reaction.
• Products can form reactants (switching).
• Dynamic equilibrium occurs in a closed area.
• Saturated Solution Example: NaCl (s) ⇌ NaCl (aq)

pH

• Measures how acidic or basic a solution is by measuring the concentration of hydrogen ions or hydronium ions [H3O+].
• Determined by the dissociation of H₂O: H₂O ⇌ H⁺ + OH⁻ or 2H₂O ⇌ H₃O⁺ + OH⁻.
• More H⁺ and H₃O⁺ than OH⁻ indicates an acidic solution.
• More OH⁻ than H₃O⁺ indicates a basic solution.
• Scale ranges from 0-14.
  • 0-6 acidic
  • 7 neutral
  • 8-14 alkaline/basic

Calculating pH

• Strong Acids: pH = -log[H+], where [H+] is the molarity of hydrogen ions.
• Strong Bases: pOH = -log[OH-], pH = 14 - pOH, where [OH-] is the molarity of hydroxide ions.
• The pH of a solution is determined by the concentration of H⁺ or H₃O⁺ ions.

Arrhenius Model

• Acids produce hydrogen ions in aqueous solutions.
• Aqueous means the solvent is water.
• Bases produce OH⁻ in aqueous solutions.
• Example: HCl → H⁺ + Cl⁻
• Acid: forms H⁺ in water.
• Example: NaOH → Na⁺ + OH⁻
• Base: forms OH⁻ in water.

Bronsted-Lowry Model

• Acid: proton donor [H+].
• Base: proton acceptor.
• Whatever is lost by the acid is gained by the base
• Bronsted acids dissociate to increase H⁺ concentration in solution.
• Bronsted bases dissociate by taking a proton to produce OH⁻.
• Example: HCl + NH₃ → NH₄⁺ + Cl⁻ (H⁺ donates a proton to form Cl⁻).
• Conjugate Acid: formed when a base accepts a proton.
• Conjugate Base: formed when an acid donates a proton.

Enthalpy of Solutions

• The formation of a solution can be spontaneous, but not guaranteed.
• Positive Enthalpy: endothermic system (enthalpy > 0).
• Negative Enthalpy: exothermic system (losing potential energy, dissolving solvent releases energy, e<1).
• Breaking solute-solute IMFs requires energy (endothermic).
• Forming solute-solvent interaction bonds is exothermic (releasing energy).
• Enthalpy: how much energy is possible in a solution.
• Net solution energy (positive or negative) = total energy shift between solute-solute and solute-solvent forces (calculate difference).

Solvent (Water)

• Particles in H₂O are free moving and have kinetic energy.
• H₂O ⇌ H⁺ + OH⁻
• The energy from the system makes it endothermic.
• They will ionize and will be rereleased again into the solute breaking the bonds (take it in and release)
• Molecular equilibrium in terms of energy:
  • The enthalpy for breaking the bonds is more than the solution (exothermic).
  • ΔH₁ + ΔH₂ > ΔH₃

Strength of Acids/Bases

• Refers to how completely an acid or base ionizes in solution.
• Strong acids completely ionize in solution.
• Weak acids partially dissociate in solution, establishing equilibrium.
• Strong bases completely ionize in solution.
• Weak bases partially ionize in solution, establishing equilibrium.
• A weak acid does not give all of its hydrogen atoms.
• For a strong acid, the arrow goes one way, indicating complete ionization of hydrogen.
• Acids with two arrows only lose one hydrogen ion.
• For a weak acid with two arrows, it partially gives the hydrogen.