Properties of Solutions

Introduction to Solutions

Solutions are homogeneous mixtures of two or more pure substances, where the solute is dispersed uniformly throughout the solvent. The ability of substances to form solutions depends on the natural tendency toward mixing and the strength of intermolecular forces.

• Homogeneous mixture: Uniform composition throughout.

• Solute: Substance dissolved in the solvent.

• Solvent: Substance in which the solute is dissolved.

Natural Tendency Toward Mixing (13.1)

Mixing of gases is a spontaneous process, increasing the randomness (entropy) of the system. The formation of solutions is favored by the increase in entropy that accompanies mixing.

• Entropy: Thermodynamic quantity representing disorder or randomness.

• Mixing increases entropy, making solution formation spontaneous.

Intermolecular Forces of Attraction

Intermolecular forces between solute and solvent molecules determine the ability to form solutions. Types include dispersion forces, dipole-dipole interactions, hydrogen bonding, and ion-dipole interactions.

• Dispersion forces: Present in all molecules, especially nonpolar.

• Dipole-dipole: Between polar molecules.

• Hydrogen bonding: Strong dipole-dipole involving H and N, O, or F.

• Ion-dipole: Between ions and polar molecules.

Attractions Involved When Forming a Solution

Three types of interactions must be considered:

• Solute–solute interactions: Must be overcome to disperse solute particles.

• Solvent–solvent interactions: Must be overcome to make room for solute.

• Solvent–solute interactions: Occur as particles mix.

Solvation (Hydration)

Solvation is the process by which solvent molecules surround and interact with solute ions or molecules. When water is the solvent, this process is called hydration.

• Hydration: Water molecules surround ions, stabilizing them in solution.

Energetics of Solution Formation

The formation of a solution involves changes in enthalpy. For an endothermic process, the increase in entropy can drive solution formation. Exothermic solutions are always spontaneous.

• ΔHmix: Enthalpy change of mixing.

• Spontaneity depends on both enthalpy and entropy.

Saturated Solutions and Solubility (13.2)

Saturated, Unsaturated, and Supersaturated Solutions

The solution-making process and crystallization are opposing processes. When the rates are equal, the solution is saturated. If less solute is present, the solution is unsaturated. Supersaturated solutions contain more solute than is normally possible and are unstable.

• Saturated solution: Maximum solute dissolved; equilibrium with undissolved solute.

• Unsaturated solution: Less than maximum solute dissolved.

• Supersaturated solution: More solute than equilibrium; unstable.

Solubility

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.

• Solubility: Maximum solute dissolved at a given temperature.

Factors That Affect Solubility (13.3)

Solute–Solvent Interactions

"Like dissolves like" is a guiding principle. Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes. The strength of solute–solvent interactions determines solubility.

• Stronger interactions increase solubility.

Organic Molecules in Water

Polar organic molecules dissolve better in water than nonpolar ones. Hydrogen bonding increases solubility.

• Hydrogen bonding sites enhance solubility.

Liquid/Liquid Solubility

Liquids that mix in all proportions are miscible; those that do not are immiscible. Polar and nonpolar liquids are typically immiscible.

• Miscible: Mix in all proportions.

• Immiscible: Do not mix.

Solubility and Biological Importance

Fat-soluble vitamins are nonpolar and stored in fatty tissue. Water-soluble vitamins must be consumed regularly.

• Vitamin A: Fat-soluble, nonpolar.

• Vitamin C: Water-soluble, polar.

Predicting Solubility Patterns

Solubility depends on molecular structure and polarity. Nonpolar solutes dissolve in nonpolar solvents; ionic and polar solutes dissolve in polar solvents.

• C7H16, I2: Nonpolar, soluble in CCl4.

• Na2SO4, HCl: Ionic/polar, soluble in water.

Pressure Effects

The solubility of gases increases with pressure, described by Henry's Law. Solubility of solids and liquids is not appreciably affected by pressure.

• Henry's Law:

Temperature Effects

For most solids, solubility increases with temperature. For gases, solubility decreases as temperature increases.

• Cold water holds more dissolved gas than warm water.

Henry's Law Calculation Example

Henry's Law relates the solubility of a gas to its partial pressure above the solution.

Solution Concentration (13.4)

Units of Concentration

Several units are used to express concentration:

• Mass percentage:

• Parts per million (ppm):

• Parts per billion (ppb):

• Mole fraction (χ):

• Molarity (M):

• Molality (m):

Converting Units

Dimensional analysis is used to convert between units. Density is required to convert between molality and molarity.

Sample Calculations

• Mass %:

• ppm:

Mole Fraction and Molality Calculation

• Mole fraction:

• Molality:

Molarity Using Density

Molarity is calculated using the mass of solute, mass of solvent, and density of the solution.

• Molarity:

Colligative Properties (13.5)

Colligative Properties

Colligative properties depend only on the number of solute particles, not their identity. These include vapor-pressure lowering, boiling-point elevation, freezing-point depression, and osmotic pressure.

• Vapor-pressure lowering

• Boiling-point elevation

• Freezing-point depression

• Osmotic pressure

Vapor Pressure

Nonvolatile solutes lower the vapor pressure of a solution compared to the pure solvent.

Raoult’s Law

Raoult’s Law states that the vapor pressure of a volatile solvent over a solution is the product of the mole fraction of the solvent and the vapor pressure of the pure solvent.

Boiling-Point Elevation

Lowered vapor pressure means a higher temperature is needed to reach atmospheric pressure, raising the boiling point.

Freezing-Point Depression

The freezing point of a solution is lower than that of the pure solvent.

Boiling-Point Elevation and Freezing-Point Depression Calculations

The change in temperature is proportional to molality and the van’t Hoff factor.

The van’t Hoff Factor (i)

The van’t Hoff factor accounts for dissociation in solution. For NaCl, i = 2. Actual values depend on concentration.

Osmosis and Osmotic Pressure

Osmosis is the net movement of solvent from low to high solute concentration across a semipermeable membrane. Osmotic pressure is the pressure required to stop osmosis.

Types of Solutions and Osmosis

• Isotonic: Same osmotic pressure.

• Hypotonic: Lower osmotic pressure.

• Hypertonic: Higher osmotic pressure.

Osmosis and Blood Cells

Red blood cells in hypertonic solutions shrivel (crenation); in hypotonic solutions, they burst (hemolysis). IV solutions must be isotonic.

Sample Calculation: Vapor Pressure of a Solution

• Calculate mole fraction and use Raoult’s Law.

Sample Calculation: Boiling-Point Elevation and Freezing-Point Depression

• Use molality and constants to calculate temperature changes.

Sample Calculation: Osmotic Pressure

Sample Calculation: Molar Mass from Freezing-Point Depression or Boiling-Point Elevation

• Use temperature change and mass to calculate molar mass.

Sample Calculation: Molar Mass from Osmotic Pressure

• Use osmotic pressure and mass to calculate molar mass.

Colloids (13.6)

Colloids

Colloids are suspensions of particles larger than individual ions or molecules but too small to settle out by gravity. They form the dividing line between solutions and suspensions.

• Colloid: Intermediate particle size; stable suspension.

Tyndall Effect

Colloidal suspensions scatter light, a phenomenon known as the Tyndall effect. Solutions do not scatter light.

Colloids and Biomolecules

Large molecules forming colloids in water often have hydrophilic portions facing outward to interact with water.

Stabilizing Colloids by Adsorption

Ions can adhere to the surface of hydrophobic colloids, allowing interaction with aqueous solution and stabilizing the colloid.

Colloids in Biological Systems

Colloids aid in emulsification of fats and oils in aqueous solutions. An emulsifier enables substances that normally do not dissolve to do so.

Brownian Motion

Colloidal particles exhibit Brownian motion due to collisions with smaller solvent molecules.