CHEM 120 Week 1

Nonelectrolyte

Dissolves but does not dissociate

Strong electrolyte

Dissolves and completely dissociates

Weak electrolyte

Dissolves and partially dissociates

Electrical conductivity of nonelectrolytes

Does not conduct electricity

Electrical conductivity of weak electrolytes

Conducts poorly

Electrical conductivity of strong electrolytes

Conducts very well

Reason conductivity differs

Number of ions present in solution

Soluble molecular compounds

Nonelectrolytes

Examples of nonelectrolytes

Sugar, alcohols

Strong acids

Acids that completely ionize in water

Examples of strong acids

HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, HClO₄

Strong bases

Bases that completely dissociate in water

Examples of strong bases

LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂

Soluble ionic compounds

Strong electrolytes

Salts of Group 1 metals

Strong electrolytes

Ammonium salts

Strong electrolytes

Nitrates

Strong electrolytes

Weak acids

Acids that partially ionize in water

Examples of weak acids

HF, HNO₂, H₂SO₃, H₂CO₃, H₃PO₄, CH₃COOH

Weak bases

Bases that partially accept protons

Most weak bases

Amines

Examples of weak bases

NH₃, CH₃NH₂

Energy change when solute separates

Endothermic

Energy change when solvent separates

Endothermic

Energy change when solute and solvent mix

Exothermic

Heat of solution (ΔHₛₒₗₙ)

Total enthalpy change when solution forms

ΔHₛₒₗₙ equation

ΔHsolute + ΔHsolvent + ΔHmix

Exothermic solution

ΔHₛₒₗₙ < 0

Endothermic solution

ΔHₛₒₗₙ > 0

Lattice energy

Energy required to separate ionic solid into gaseous ions

Sign of lattice energy

Always positive

Solvation

Process of solvent surrounding solute particles

Hydration

Solvation in water

Sign of hydration energy

Always negative

ΔHₛₒₗₙ for ionic compounds

ΔHlattice + ΔHhydration

Entropy

Measure of disorder or freedom of motion

State with highest entropy

Gas

Entropy change when forming solutions

Usually increases

Saturated solution

Contains maximum dissolved solute at given temperature

Unsaturated solution

Contains less than maximum solute

Supersaturated solution

Contains more solute than equilibrium allows

Effect of temperature on ionic solubility

Usually increases solubility

Solubility of gases in water

Low

Effect of temperature on gas solubility

Decreases as temperature increases

Henry's Law

Gas solubility is proportional to gas pressure

Henry's Law equation

S = kH × P

Freezing point depression

Solution freezes at lower temperature than pure solvent

Freezing point depression equation

ΔTf = Kf · m

Boiling point elevation

Solution boils at higher temperature than pure solvent

Boiling point elevation equation

ΔTb = Kb · m

Colligative properties depend on

Number of solute particles

van't Hoff factor (i)

Number of particles produced per formula unit

i for nonelectrolytes

1

i for strong electrolytes

Number of ions formed

Ion pairing

Reduces actual van't Hoff factor

Raoult's Law

Vapor pressure proportional to mole fraction of solvent

Vapor pressure of solution vs pure solvent

Always lower

Osmosis

Flow of solvent through semipermeable membrane

Osmotic pressure

Pressure needed to stop osmosis

Osmotic pressure equation

Π = MRT

Isotonic solution

Same concentration as cell

Hypertonic solution

Higher concentration than cell

Hypotonic solution

Lower concentration than cell

Colloid

Heterogeneous mixture with dispersed particles

Tyndall effect

Light scattering by colloids