Properties of Solutions – Chapter 13 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 ability of substances to form solutions depends on the natural tendency toward mixing and the strength of intermolecular forces.

• Solute: The substance being dissolved.

• Solvent: The substance doing the dissolving, usually present in greater amount.

• Homogeneous mixture: A mixture with uniform composition throughout.

Natural Tendency Toward Mixing

Mixing of gases and liquids is a spontaneous process that increases the randomness, or entropy, of the system. The formation of solutions is favored by the increase in entropy that accompanies mixing.

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

• Spontaneous mixing increases entropy.

Intermolecular Forces of Attraction

Intermolecular forces are responsible for the interactions between solute and solvent molecules. These forces must be considered when forming a solution:

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

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

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

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.

• Solvation stabilizes solute particles in solution.

• Hydration is crucial for dissolving ionic compounds in water.

Energetics of Solution Formation

The formation of a solution involves energetic considerations. For an endothermic process to occur, the increase in entropy must compensate for the energy required. Exothermic solution formation is always spontaneous.

• Endothermic process: Absorbs energy; may occur if entropy increases sufficiently.

• Exothermic process: Releases energy; generally spontaneous.

Aqueous Solution Versus Chemical Reaction

Not all substances that disappear in a solvent are dissolved; some may react chemically. For example, nickel reacts with hydrochloric acid rather than simply dissolving.

• Dissolution: Physical process of solute dispersing in solvent.

• Chemical reaction: Formation of new substances.

Opposing Processes: Solution Formation and Crystallization

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

• Saturated solution: Contains the maximum amount of dissolved solute.

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

Solubility

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature. Saturated solutions have this amount dissolved; unsaturated solutions have less.

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

Factors Affecting Solubility

Several factors influence solubility:

• Solute–solvent interactions: Stronger interactions increase solubility.

• Pressure: Affects solubility of gases (see Henry’s Law).

• Temperature: Generally increases solubility of solids, decreases solubility of gases.

Solute–Solvent Interactions: "Like Dissolves Like"

Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes. Hydrogen bonding increases solubility in water.

• Miscible liquids: Mix in all proportions.

• Immiscible liquids: Do not mix (e.g., hexane and water).

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.

Pressure Effects and Henry’s Law

Pressure has little effect on the solubility of solids and liquids, but greatly affects gases. According to Henry’s Law:

• The solubility of a gas is proportional to its partial pressure above the solution.

$S_g = k_H \cdot P_g$

• $S_g$ = solubility of the gas

• $k_H$ = Henry's law constant

• $P_g$ = partial pressure of the gas

Temperature Effects

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

• Cold water holds more dissolved oxygen than warm water.

Solution Concentration

Concentration expresses the amount of solute in a given amount of solution. Common units include:

• Mass percentage

• Parts per million (ppm)

• Parts per billion (ppb)

• Mole fraction

• Molarity (M)

• Molality (m)

Units of Concentration

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

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

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

• Mole fraction ($\chi$): $\chi = \frac{\text{moles of component}}{\text{total moles in solution}}$

• 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 Versus Molality

When water is the solvent, dilute solutions have similar molarity and molality. Molality does not vary with temperature, while molarity does (since volume changes with temperature).

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

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

• Osmotic pressure

Vapor Pressure and Raoult’s Law

Adding a nonvolatile solute lowers the vapor pressure of the solvent. Raoult’s Law states:

$P_{solution} = \chi_{solvent} \cdot P^0_{solvent}$

• $P_{solution}$ = vapor pressure of the solution

• $\chi_{solvent}$ = mole fraction of the solvent

• $P^0_{solvent}$ = vapor pressure of pure solvent

Boiling-Point Elevation and Freezing-Point Depression

Lowered vapor pressure means a higher boiling point and a lower freezing point. The change in temperature is proportional to molality and the van’t Hoff factor ($i$):

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

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

• $\Delta T_b$ = boiling-point elevation

• $\Delta T_f$ = freezing-point depression

• $K_b, K_f$ = constants for solvent

• $m$ = molality

• $i$ = van’t Hoff factor (number of particles formed when solute dissolves)

Osmosis and Osmotic Pressure

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

$\Pi = i \cdot M \cdot R \cdot T$

• $\Pi$ = osmotic pressure

• $i$ = van’t Hoff factor

• $M$ = molarity

• $R$ = gas constant

• $T$ = temperature (K)

Types of Solutions and Osmosis

• Isotonic: Same osmotic pressure; solvent passes at equal rates.

• Hypotonic: Lower osmotic pressure; solvent leaves faster than it enters.

• Hypertonic: Higher osmotic pressure; solvent enters faster than it leaves.

Osmosis and Blood Cells

Red blood cells have semipermeable membranes. In hypertonic solutions, they shrivel (crenation); in hypotonic solutions, they burst (hemolysis). Intravenous solutions must be isotonic to blood.

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.

• Colloids scatter light (Tyndall effect).

• Colloids can be stabilized by adsorption of ions.

• Colloids aid in emulsification of fats and oils in biological systems.

Tyndall Effect

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

Colloids and Biomolecules

Large molecules forming colloids in water often have hydrophilic (water-loving) ends facing outward and hydrophobic (water-fearing) ends inward. This structure aids in stability and interaction with water.

Stabilizing Colloids by Adsorption

Ions can adhere to the surface of hydrophobic colloids, allowing them to interact with aqueous solutions and remain dispersed.

Colloids in Biological Systems

Colloids help emulsify fats and oils in aqueous solutions. Emulsifiers enable 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, keeping them suspended.

*Additional info: Some tables referenced (e.g., solubilities of gases, types of colloids) were not fully visible; key concepts and definitions have been inferred and expanded for completeness.*