Chapter 13: Properties of Solutions – General Chemistry Study Notes

Properties of Solutions

Introduction to Solutions

Solutions are homogeneous mixtures composed of two or more pure substances. In a solution, the solute is uniformly dispersed throughout the solvent. The formation of solutions depends on the natural tendency toward mixing and the nature of intermolecular forces.

• Solute: The substance dissolved in the solvent.

• Solvent: The substance present in the greatest amount, which dissolves the solute.

• Homogeneous mixture: Uniform composition throughout.

Natural Tendency Toward Mixing and Entropy Mixing of gases is a spontaneous process that increases the randomness (entropy) of the system. The tendency toward greater entropy favors the formation of solutions.

• Entropy (S): A thermodynamic quantity representing the degree of disorder or randomness in a system.

• Mixing increases entropy, making solution formation more likely.

Intermolecular Forces in Solution Formation

Intermolecular forces play a crucial role in the formation of solutions. The types of forces include dispersion, dipole-dipole, hydrogen bonding, and ion-dipole interactions.

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

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

• Solvent–solute interactions: Occur as particles mix, stabilizing the solution.

Solvation and 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.

• Solvent molecules stabilize ions in solution by surrounding them.

• Hydration is especially important for ionic compounds in water.

Energetics of Solution Formation

The formation of a solution involves changes in enthalpy (ΔH) and entropy (ΔS). The overall enthalpy change (ΔHsolution) is the sum of the energies required to separate solute and solvent particles and the energy released when solute and solvent interact.

• Exothermic process: ΔHsolution < 0; solution formation releases heat and is usually spontaneous.

• Endothermic process: ΔHsolution > 0; solution formation absorbs heat, but may still be spontaneous if entropy increases sufficiently.

Aqueous Solution Versus Chemical Reaction

Not all substances that disappear in a solvent are simply dissolving; some may undergo a chemical reaction with the solvent.

• Example: Nickel metal reacts with hydrochloric acid to form NiCl2(aq) and H2(g), not just a physical dissolution.

Solubility and Types of Solutions

Opposing Processes: Dissolution and Crystallization

Dissolution and crystallization are opposing processes. When their rates are equal, the solution is saturated. If less solute is present than the maximum, the solution is unsaturated.

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

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

Supersaturated Solutions

A supersaturated solution contains more solute than is normally possible at a given temperature. These solutions are unstable and can crystallize rapidly if disturbed.

• Crystallization can be triggered by adding a seed crystal or scratching the container.

Factors Affecting Solubility

Solubility is influenced by solute–solvent interactions, pressure (for gases), and temperature.

• Solute–solvent interactions: "Like dissolves like"—polar solutes dissolve in polar solvents, nonpolar in nonpolar.

• Pressure: Affects solubility of gases only.

• Temperature: Affects solubility of both solids and gases, but in different ways.

Solute–Solvent Interactions and Examples

The stronger the interaction between solute and solvent, the greater the solubility. For gases, larger molecules with greater dispersion forces are more soluble in water.

Gas	Molar Mass (g/mol)	Solubility (M)
N2	28.0	0.69 × 10−3
O2	32.0	1.38 × 10−3
Ar	39.9	1.50 × 10−3
Kr	83.8	2.79 × 10−3

Organic Molecules in Water

Polar organic molecules dissolve better in water than nonpolar ones. Hydrogen bonding increases solubility, as C–C and C–H bonds are not very polar.

Liquid/Liquid Solubility

Liquids that mix in all proportions are miscible. Liquids that do not mix are immiscible. For example, hexane (nonpolar) and water (polar) are immiscible.

Solubility and Biological Importance

Fat-soluble vitamins (e.g., vitamin A) are nonpolar and stored in fatty tissue. Water-soluble vitamins (e.g., vitamin C) must be consumed regularly as they are not stored in the body.

Effects of Pressure and Temperature on Solubility

Pressure Effects

Pressure has little effect on the solubility of solids and liquids, but greatly affects gases. Increasing the pressure of a gas above a solution increases its solubility.

Henry’s Law

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

• Henry's Law: $S_g = kP_g$

• Where $S_g$ is the solubility, $k$ is the Henry's law constant, and $P_g$ is the partial pressure of the gas.

Temperature Effects

For most solids, solubility increases with temperature, but not always. For all gases, solubility decreases as temperature increases.

• Cold water holds more dissolved oxygen than warm water.

Concentration Units

Common Units of Concentration

• Mass percentage: $\text{Mass %} = \frac{\text{mass of component}}{\text{total mass of solution}} \times 100$

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

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

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

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

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

Molarity vs. Molality

Molarity depends on volume and changes with temperature, while molality depends on mass and does not change with temperature. For dilute aqueous solutions, values are similar.

Converting Units

To convert between molarity and molality, the density of the solution is required. Use dimensional analysis for all conversions.

Colligative Properties

Definition and Types

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

• Vapor-pressure lowering

• Boiling-point elevation

• Freezing-point depression

• Osmotic pressure

Vapor Pressure Lowering

Adding a nonvolatile solute to a solvent lowers the solvent's vapor pressure due to solute–solvent interactions.

Raoult’s Law

Raoult’s Law quantifies vapor pressure lowering:

• $P_{\text{solution}} = \chi_{\text{solvent}} P^\circ_{\text{solvent}}$

• Where $P_{\text{solution}}$ is the vapor pressure of the solution, $\chi_{\text{solvent}}$ is the mole fraction of the solvent, and $P^\circ_{\text{solvent}}$ is the vapor pressure of the pure solvent.

Boiling-Point Elevation

Since vapor pressure is lowered, a higher temperature is required to reach atmospheric pressure, raising the boiling point of the solution.

Freezing-Point Depression

The freezing point of a solution is lower than that of the pure solvent. This is also a result of vapor pressure lowering.

Quantitative Relationships for Boiling-Point Elevation and Freezing-Point Depression

• $\Delta T_b = i K_b m$

• $\Delta T_f = i K_f m$

• Where $i$ is the van’t Hoff factor (number of particles the solute forms in solution), $K_b$ and $K_f$ are the boiling-point elevation and freezing-point depression constants, and $m$ is the molality.

Solvent	Normal Boiling Point (°C)	Kb (°C/m)	Normal Freezing Point (°C)	Kf (°C/m)
Water (H2O)	100.0	0.51	0.0	1.86
Benzene (C6H6)	80.1	2.53	5.5	5.12
Ethanol (C2H5OH)	78.4	1.22	−114.6	1.99
Carbon tetrachloride (CCl4)	76.8	5.02	−22.3	29.8
Chloroform (CHCl3)	61.2	3.63	−63.5	4.68

Osmosis and Osmotic Pressure

Osmosis

Osmosis is the movement of solvent molecules through a semipermeable membrane from a region of low solute concentration to high solute concentration. The pressure required to stop this flow is the osmotic pressure.

Osmotic Pressure Equation

• $\Pi = iMRT$

• Where $\Pi$ is the osmotic pressure, $i$ is the van’t Hoff factor, $M$ is molarity, $R$ is the gas constant, and $T$ is temperature in Kelvin.

Types of Solutions and 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.

Osmosis and Blood Cells

Red blood cells have semipermeable membranes. In hypertonic solutions, they shrink (crenation); in hypotonic solutions, they swell and may burst (hemolysis). IV solutions must be isotonic to blood.

Colloids

Definition and Types

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

Phase of Colloid	Dispersing Substance	Dispersed Substance	Colloid Type	Example
Gas	Gas	—	None (all are solutions)	—
Gas	Liquid	Aerosol	Fog
Gas	Solid	Aerosol	Smoke
Liquid	Gas	Foam	Whipped cream
Liquid	Liquid	Emulsion	Milk
Liquid	Solid	Sol	Paint
Solid	Gas	Solid foam	Marshmallow
Solid	Liquid	Solid emulsion	Butter
Solid	Solid	Solid sol	Ruby glass

Tyndall Effect

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

Colloids and Biomolecules

Many biomolecules have both hydrophilic (water-loving) and hydrophobic (water-fearing) regions. In water, hydrophilic portions face outward, stabilizing the colloid.

Stabilizing Colloids by Adsorption

Ions can adsorb onto the surface of hydrophobic colloids, allowing them to interact with water and remain dispersed.

Colloids in Biological Systems

Colloids help emulsify fats and oils in aqueous solutions. An emulsifier enables substances that normally do not dissolve in a solvent to do so.

Brownian Motion

Colloidal particles exhibit Brownian motion due to collisions with much smaller solvent molecules. This motion helps keep colloidal particles suspended.

Radius of Sphere (nm)	Mean Free Path (mm)
1	1.23
10	0.390
100	0.123
1000	0.039

Additional info: These notes cover all major aspects of solution properties, including types of solutions, factors affecting solubility, colligative properties, and the nature and behavior of colloids, as relevant to a General Chemistry curriculum.