Chapter 13: Properties of Solutions – General Chemistry Study Notes

Properties of Solutions

Introduction to Solutions

Solutions are a fundamental concept in chemistry, representing homogeneous mixtures of two or more pure substances. The solute is the substance dispersed uniformly throughout the solvent. The ability of substances to form solutions depends on both the natural tendency toward mixing and the intermolecular forces present.

• Solution: Homogeneous mixture of two or more substances.

• Solute: Substance present in a lesser amount, dispersed in the solvent.

• Solvent: Substance present in a greater amount, dissolves the solute.

• Key factors: Natural tendency toward mixing (entropy) and intermolecular forces.

Natural Tendency Toward Mixing (Section 13.1)

Entropy and Spontaneity

Mixing of gases and formation of solutions are spontaneous processes driven by an increase in entropy, a thermodynamic quantity representing randomness or disorder.

• Mixing increases randomness (entropy) in the system.

• Formation of solutions is favored by the increase in entropy that accompanies mixing.

• Each gas in a mixture behaves independently, filling the container as if it were alone.

Intermolecular Forces of Attraction

Types of Intermolecular Forces

Intermolecular forces are crucial in determining whether substances will form solutions. These forces can occur between solute and solvent molecules and include:

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

• Dipole-dipole interactions: Occur between polar molecules.

• Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

• Ion-dipole interactions: Occur between ions and polar molecules (e.g., Na+ in water).

Energetics of Solution Formation

Enthalpy Changes

The formation of a solution involves overcoming solute-solute and solvent-solvent interactions, and forming solute-solvent interactions. The overall enthalpy change ($\Delta H_{\text{solution}}$) determines whether the process is endothermic or exothermic.

• Endothermic process: $\Delta H_{\text{mix}}$ must be close to the sum of $\Delta H_{\text{solute}}$ and $\Delta H_{\text{solvent}}$; entropy becomes important.

• Exothermic process: Spontaneous if $\Delta H_{\text{solution}}$ is negative.

Saturated Solutions and Solubility (Section 13.2)

Dynamic Equilibrium and Types of Solutions

The process of dissolving and crystallization are opposing processes. When their rates are equal, the solution is saturated.

• Saturated solution: Contains the maximum amount of solute that can dissolve at a given temperature.

• Unsaturated solution: Contains less solute than the maximum amount possible.

• Supersaturated solution: Contains more solute than is normally possible at that temperature; unstable and can crystallize easily.

Factors Affecting Solubility (Section 13.3)

Solute-Solvent Interactions

• "Like dissolves like": Polar solutes dissolve in polar solvents; nonpolar in nonpolar.

• Stronger solute-solvent interactions increase solubility.

• For gases, larger molecules with greater dispersion forces are more soluble in water.

Pressure Effects

• Solubility of solids and liquids is not significantly affected by pressure.

• Gas solubility increases with increasing pressure (see Henry's Law).

Temperature Effects

• For most solids, solubility increases with temperature (exceptions exist).

• For gases, solubility decreases as temperature increases.

Henry's Law

The solubility of a gas in a liquid is proportional to the partial pressure of the gas above the solution.

• Equation: $S_g = k P_g$

• $S_g$ = solubility of the gas, $k$ = Henry's law constant, $P_g$ = partial pressure of the gas

Solution Concentration (Section 13.4)

Units of Concentration

• Mass percentage: $\frac{\text{mass of component in soln}}{\text{total mass of soln}} \times 100$

• Parts per million (ppm): $\frac{\text{mass of component in soln}}{\text{total mass of soln}} \times 10^6$

• Parts per billion (ppb): $\frac{\text{mass of component in soln}}{\text{total mass of soln}} \times 10^9$

• Mole fraction ($x$): $\frac{\text{moles of component}}{\text{total moles of all components}}$

• Molarity (M): $\frac{\text{moles of solute}}{\text{liters of soln}}$

• Molality (m): $\frac{\text{moles of solute}}{\text{kilograms of solvent}}$

Molarity vs. Molality

• Molarity depends on volume (changes with temperature).

• Molality depends on mass (does not change with temperature).

Converting Units

• To convert between molality and molarity, the density of the solution is required.

Colligative Properties (Section 13.5)

Definition and Types

Colligative properties depend only on the number of solute particles, not their identity. These include:

• Vapor-pressure lowering

• Boiling-point elevation

• Freezing-point depression

• Osmotic pressure

Raoult's Law

• The vapor pressure of a solution is the product of the mole fraction of the solvent and the vapor pressure of the pure solvent.

• Equation: $P_{\text{solution}} = X_{\text{solvent}} P_{\text{solvent}}$

Boiling-Point Elevation and Freezing-Point Depression

• Both are proportional to the molality of the solution and the van't Hoff factor ($i$).

• Boiling-point elevation: $\Delta T_b = i K_b m$

• Freezing-point depression: $\Delta T_f = -i K_f m$

• $K_b$ and $K_f$ are constants for the solvent.

Osmosis and Osmotic Pressure

• Osmosis: Net movement of solvent from low to high solute concentration across a semipermeable membrane.

• Osmotic pressure ($\Pi$): $\Pi = iMRT$

• $M$ = molarity, $R$ = gas constant, $T$ = temperature in Kelvin, $i$ = van't Hoff factor

Types of Solutions in Osmosis

• Isotonic: Same osmotic pressure; no net movement of solvent.

• Hypotonic: Lower osmotic pressure; solvent leaves this solution.

• Hypertonic: Higher osmotic pressure; solvent enters this solution.

Colloids (Section 13.6)

Definition and Properties

Colloids are mixtures with particle sizes between those of solutions and suspensions. They do not settle out by gravity and exhibit unique properties such as the Tyndall effect (scattering of light).

• Colloids: Particles 10–1000 nm in size, intermediate between solutions and suspensions.

• Tyndall effect: Scattering of light by colloidal particles.

• Stabilization: Ions or molecules can adsorb to colloid surfaces, stabilizing them in solution.

• Biological importance: Colloids aid in emulsification of fats and oils in aqueous solutions.

Property	Solutions	Colloids	Suspensions
Appearance	Clear, homogeneous	Cloudy, homogeneous	Cloudy, heterogeneous
Particle Size	Molecule size	10–1000 nm	>1000 nm
Tyndall Effect	No	Yes	Yes
Sedimentation	No	No	Yes

Brownian Motion

• Random motion of colloidal particles due to collisions with solvent molecules.

Summary Table: Key Equations

Property	Equation
Mass %	$\frac{\text{mass of component}}{\text{total mass of solution}} \times 100$
ppm	$\frac{\text{mass of component}}{\text{total mass of solution}} \times 10^6$
ppb	$\frac{\text{mass of component}}{\text{total mass of solution}} \times 10^9$
Mole Fraction	$\frac{\text{moles of component}}{\text{total moles}}$
Molarity (M)	$\frac{\text{moles of solute}}{\text{liters of solution}}$
Molality (m)	$\frac{\text{moles of solute}}{\text{kg of solvent}}$
Henry's Law	$S_g = k P_g$
Raoult's Law	$P_{\text{solution}} = X_{\text{solvent}} P_{\text{solvent}}$
Boiling-Point Elevation	$\Delta T_b = i K_b m$
Freezing-Point Depression	$\Delta T_f = -i K_f m$
Osmotic Pressure	$\Pi = iMRT$

Additional info: This summary includes expanded explanations, definitions, and context for all major points covered in the provided lecture slides, suitable for a General Chemistry college course.