Physical Properties of Solutions

Physical Properties of Solutions

General Overview

• Course: General Chemistry II - CHEM 2061

• Semester: Spring 2026

• Instructor: Dr. Niki Shoup

Types of Solutions

• Chemical interactions typically occur between two compounds in a solvent.

• Table 12.1 - Types of Solutions by State:

  • Gas & Gas → Resulting Solution: Gas (e.g., Air)

  • Gas in Liquid → Resulting Solution: Gas (e.g., Soda water - CO2 in water)

  • Gas in Solid → Resulting Solution: Solid (e.g., H2 gas in palladium)

  • Liquid in Liquid → Resulting Solution: Liquid (e.g., Ethanol in water)

  • Solid in Liquid → Resulting Solution: Liquid (e.g., NaCl in water)

  • Liquid in Solid → Resulting Solution: Solid (e.g., Brass - Cu/Zn)

  • Solid in Solid → Resulting Solution: Solid (e.g., Solder - Sn/Pb)

Types of Solutions Defined

• Saturated Solutions: Contains the maximum amount of solute that can dissolve in a solvent.

• Solute: The substance that is dissolved in a solution.

• Solvent: The medium in which the solute is dissolved.

• Unsaturated Solutions: Contains less than the maximum amount of solute that can dissolve in a solvent.

• Supersaturated Solutions: Contains more solute than is possible under normal circumstances; excess solute will precipitate out.

• Crystallization: The process of solute particles forming solid crystals as they precipitate out of solution.

Solutions on the Molecular Level

• Solvent-Solvent Interactions: Interactions between solvent molecules.

• Solute-Solute Interactions: Interactions between solute molecules.

• Solute-Solvent Interactions: Interactions that occur between solute and solvent.

• Endothermic Solvation: A process requiring energy.

• Exothermic Solvation: A process that releases energy.

Miscible Liquids and Solvated Solids

• Miscible Liquids: Two liquids that fully dissolve in each other (e.g., ethanol in water).

• Non-miscible Liquids: Liquids that do not mix (e.g., oil in water).

• Solvation: The process in which a solid dissolves into a liquid.

• General principle: “Like dissolves like”; similar intermolecular forces facilitate dissolution.

Solubility Analysis

Solubility in Water vs. Benzene

• Determine which of the following compounds are soluble in water and which are soluble in benzene:

  • H2O (water)

  • Benzene

  • LiCl

  • H3C (methyl)

  • OH (hydroxyl group)

Quantitative Description of Solubility

• Percent by Mass (Unitless):
  ext{Mass percent} = rac{ ext{mass of solute}}{ ext{total mass of solution}} imes 100 ext{%}

• Mole Fraction (Unitless):
  X_A = rac{ ext{moles of compound A}}{ ext{total moles of compounds in solution}}

• Molarity (mol/L):
  M = rac{ ext{moles of solute}}{ ext{liters of solvent}}

• Molality (mol/kg):
  m = rac{ ext{moles of solute}}{ ext{mass of solvent in kg}}

Application of Units for Solubility Calculation

• Percent by Mass: Used when only the mass of the solute and mass of the solution are known.

• Mole Fraction: Useful for gas phase calculations and vapor pressure determinations.

• Molarity: Commonly used due to ease of measurement; applicable for solids and liquids.

• Molality: Better for solutions when temperature variations may affect volume measurements.

Practical Examples in Solubility Calculations

Example Problem

• Density of an aqueous solution containing 10.0% ethanol (C2H5OH) by mass is 0.984 g/mL.

  • (a) Calculate the molality of this solution.

  • (b) Calculate its molarity.

  • (c) Determine the volume of the solution that would contain 0.125 moles of ethanol.

Temperature and Solubility

• For most solids, solubility increases with temperature; used for fractional crystallization, especially with compounds of differing solubilities.

• For gases, increased temperatures lead to decreased solubility in liquids (e.g., O2 solubility in water decreases).

Pressure and Gas Solubility

• Henry's Law: Describes the relationship between gas solubility and pressure. c = kP

  • Where:

  • c = Molar concentration of gas in solution (mol/L)

  • P = Pressure of gas (atm)

  • k = Temperature-dependent constant (mol/L·atm)

• Increased pressure results in increased gas solubility.

Henry's Law Practice Example

• Solubility of nitrogen gas at 25°C and 1 atm is 6.8 imes 10^{-4} mol/L. Calculate the concentration (molarity) of nitrogen dissolved in water under atmospheric conditions with a partial nitrogen pressure of 0.78 atm.

Colligative Properties of Nonelectrolyte Solutions

• Colligative Properties: Properties that depend on the number of solute particles, not their identity.

  • Generally apply to dilute solutions (< 0.2 M).

• Examples include:

  • Vapor-pressure lowering

  • Boiling-point elevation

  • Freezing-point depression

  • Osmotic pressure

Vapor-Pressure Lowering

• Nonvolatile solute lowers vapor pressure compared to pure solvent.

• Raoult’s Law: P1 = X1 P_1°

  • Where:

  • P_1 = vapor pressure of solvent over solution

  • X_1 = mole fraction of solvent

  • P_1° = vapor pressure of pure solvent.

• For a single solute, can reformulate as:
  P1° - P1 = riangle P = X2 P1°

Raoult’s Law Practice Example

• Calculate the vapor pressure of a solution made by dissolving 82.5 g of urea (m.m. = 60.06 g/mol) in 212 mL of water at 35 °C. The vapor pressure of water at this temperature is 42.18 mmHg.

Ideal Solutions

• Lower vapor pressure due to increased disorder in solutions; pure liquids exhibit higher vapor pressure due to disorder in vapor.

• When both components of a solution are volatile, total vapor pressure is:
  PT = XA PA° + XB P_B°

Nonideal Solutions

• Describes conditions under which ideal behavior fails (for example, in strong solute-solvent attractions or solvent-solvent interaction dominating).

Fractional Distillation

• Techniques for separation of liquids based on boiling points.

• Commonly utilized in industries for refining petroleum, isolating commodity chemicals, and producing spirits.

Boiling Point Elevation

• Presence of nonvolatile solute leads to increased boiling point: riangle Tb = Kb m

  • Where:

  • riangle T_b = change in boiling point

  • K_b = molal boiling point elevation constant

  • m = molality.

  • Molality is preferred over molarity as temperature changes affect the solution volume.

Freezing Point Depression

• Similar increase observed when a solute is added: riangle Tf = Kf m

  • Where:

  • riangle T_f = change in freezing point

  • K_f = molal freezing point depression constant.

• Practical application with salts on icy roads.

Ethylene Glycol Example

• Given 651 g of ethylene glycol (EG) in 2505 g of water, calculate freezing point. Note: Molar mass of ethylene glycol is 62.01 g.

Osmotic Pressure

• Osmosis: Selective passage of solvent molecules through a semipermeable membrane from dilute to concentrated.

• Osmotic Pressure: Pressure required to halt osmosis: ext{π} = M R T

  • Where:

  • π = osmotic pressure

  • M = molarity

  • R = gas constant

  • T = absolute temperature.

Osmotic Pressure Practice Example

• The average osmotic pressure of seawater is about 30.0 atm at 25°C. Calculate the molar concentration of an aqueous sucrose solution that is isotonic with seawater.

Using Colligative Properties to Determine Molar Mass

• A 7.85 g sample of a compound with empirical formula C5H4 is dissolved in 301 g of benzene. Freezing point depression noted is 1.05°C from pure benzene. Calculate molar mass and molecular formula of compound.

Colligative Properties of Electrolyte Solutions

• For solutions containing ions, we account for the ions positioned in the solution from electrolytes.

  • Example:SEE PDF NaCl
    ightarrow Na^+ + Cl^- and CaCl_2
    ightarrow Ca^{2+} + 2Cl^- .

• Include the van't Hoff factor i in calculations for colligative properties:

  • riangle Tb = iKbm

  • riangle Tf = iKfm

  • π = iMR T .

Examplar Calculations for Electrolyte Solutions

• Osmotic pressure of a 0.010 M potassium iodide (KI) solution at 25°C is 0.465 atm; determine the van't Hoff factor for KI.

• Freezing point depression of a 0.100 m MgSO4 solution is 0.225°C; calculate van't Hoff factor for MgSO4 .

Colloids

• Definition: Colloids are dispersions of particles throughout a dispersing medium.

  • Examples of colloid classifications include: SEE PDF
    | Dispersing Medium | Dispersed Phase | Name | Example |
    |-------------------|----------------|------------|----------------------------|
    | Gas | Liquid | Aerosol | Fog, mist |
    | Gas | Solid | Aerosol | Smoke |
    | Liquid | Gas | Foam | Whipped cream |
    | Liquid | Liquid | Emulsion | Mayonnaise |
    | Liquid | Solid | Sol | Milk of magnesia |
    | Solid | Gas | Foam | Plastic foams |
    | Solid | Liquid | Gel | Jelly, butter |
    | Solid | Solid | Solid sol | Certain alloys (steel), opal|

Water Colloids

• Hydrophilic Colloids: Solutions containing large particles (e.g., proteins), made soluble through hydrogen bonding.

• Hydrophobic Colloids: Solutions such as oil in water; can be dispersed using agents like soap.