Chapter 14: Solutions – Structure, Energetics, and Solubility

Solutions

Solution Terminology

Understanding solutions is fundamental in chemistry, as they represent homogeneous mixtures of two or more substances. The terminology associated with solutions helps classify and describe their properties and behaviors.

• Solution: A homogeneous mixture of two or more substances.

• Solvent: The component present in the greater amount; typically determines the phase of the solution.

• Solute: The component present in the lesser amount; dissolved in the solvent.

Example: In an aqueous solution of salt water, water is the solvent and salt is the solute.

Table Purpose: Classification of solution types based on the phases of solute and solvent, with examples.

Solution Phase	Solute Phase	Solvent Phase	Example
Gaseous solution	Gas	Gas	Air (mainly oxygen and nitrogen)
Liquid solution	Gas	Liquid	Club soda (CO2 and water)
Liquid solution	Liquid	Liquid	Vodka (ethanol and water)
Liquid solution	Solid	Liquid	Seawater (salt and water)
Solid solution	Solid	Solid	Brass (copper and zinc) and other alloys

Solubility

Solubility quantifies the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature. It is commonly expressed as grams of solute per 100 mL or 100 g of solvent.

• Miscible: Liquids that are soluble in each other in all proportions.

• Immiscible: Liquids that are not soluble in each other.

Saturated, Unsaturated, and Supersaturated Solutions

Solutions can be classified based on the amount of solute dissolved relative to the equilibrium solubility at a given temperature.

• Saturated Solution: Contains the maximum amount of dissolved solute at equilibrium; excess solute remains undissolved.

• Unsaturated Solution: Contains less than the equilibrium amount of solute.

• Supersaturated Solution: Contains more solute than is normally possible; unstable and prone to crystallization.

Example: If 110 g CuCl2 are added to 100 mL water at 0°C (solubility = 70.6 g/100 mL), only 70.6 g dissolve (saturated), and the rest remains undissolved.

Energetics of Making Solutions

Why Do Molecular Solutions Form?

The formation of a solution depends on the types of intermolecular forces (IMFs) involved and the entropy change during mixing.

• IMFs: The strength and type of interactions (solute-solute, solvent-solvent, solute-solvent) determine the enthalpy change.

• Entropy: Mixing increases randomness (entropy), favoring solution formation even if the process is endothermic.

Conditions That Favor Solution Formation

Solution formation is favored when solute-solvent interactions are stronger than solute-solute and solvent-solvent interactions.

• Weak solute-solute interactions

• Weak solvent-solvent interactions

• Strong solute-solvent interactions

Energy Changes During Solution Formation

The process of solution formation involves three energetically distinct steps:

1. Breaking solute-solute interactions ($\Delta H_{solute}$, endothermic)

2. Breaking solvent-solvent interactions ($\Delta H_{solvent}$, endothermic)

3. Mixing solute and solvent particles ($\Delta H_{mix}$, exothermic)

The overall enthalpy change is:

$\Delta H_{solution} = \Delta H_{solute} + \Delta H_{solvent} + \Delta H_{mix}$

Exothermic: If the energy released in mixing is greater than the energy required to separate solute and solvent, the process is exothermic.

Endothermic: If the energy required to separate solute and solvent is greater than the energy released in mixing, the process is endothermic.

Application of Solution Concepts

Several real-world examples illustrate the principles of solution formation:

• Oil and Water: Immiscible due to weak solute-solvent interactions and strong solvent-solvent interactions (hydrogen bonding).

• Carbon Tetrachloride and Benzene: Both nonpolar; similar IMFs allow miscibility.

• Heptane and Octane: Both nonpolar hydrocarbons; miscible due to similar London Dispersion Forces.

• Methanol, Ethanol, and 1,2-Ethylene Glycol: Miscible in water due to hydrogen bonding.

• Vitamin C: Polar, water-soluble.

• Vitamin A: Nonpolar, fat-soluble.

Key Principle: "Like dissolves like" – polar or ionic solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents.

Trends in Alcohol Solubility in Water

As the hydrocarbon chain in alcohols increases, the molecule becomes less polar and its solubility in water decreases.

Alcohol	Formula	Solubility (g/100 g H2O)
Methanol	CH3OH	Miscible
Ethanol	CH3CH2OH	Miscible
1-Propanol	CH3CH2CH2OH	Miscible
1-Butanol	CH3CH2CH2CH2OH	7.9
1-Pentanol	CH3CH2CH2CH2CH2OH	2.7
1-Hexanol	CH3CH2CH2CH2CH2CH2OH	0.6

Ionic Solutions and Enthalpy of Hydration

Ionic Solutions

When ionic compounds dissolve in water, ions are stabilized by hydration due to ion-dipole attractions. Each ion is surrounded by several water molecules.

Enthalpy of Hydration

The enthalpy of solution for ionic compounds is given by:

$\Delta H_{solution} = -\Delta H_{lattice} + \Delta H_{hydration}$

Lattice Energy ($\Delta H_{lattice}$): The energy released when one mole of an ionic solid forms from its gaseous ions; highly exothermic.

Enthalpy of Hydration ($\Delta H_{hydration}$): The energy released when one mole of gaseous ions is dissolved in water; always exothermic.

• The higher the charge density (charge/volume) of the ion, the more exothermic the enthalpy of hydration.

• For ionic compounds, the overall solubility depends on the balance between lattice energy and hydration energy.

Effect of Temperature and Pressure on Solubility

Temperature and pressure significantly affect solubility:

• Solids: Solubility generally increases with temperature, but there are exceptions (e.g., Na2SO4).

• Gases: Solubility decreases with increasing temperature (thermal pollution).

• Pressure: Solubility of gases increases with partial pressure, described by Henry's Law:

$S_{gas} = k_H P_{gas}$

Where $S_{gas}$ is the solubility of the gas, $k_H$ is Henry's Law constant, and $P_{gas}$ is the partial pressure.

Summary Table: Key Concepts in Solution Formation

Concept	Definition	Example/Application
Solution	Homogeneous mixture	Salt water
Solubility	Maximum solute dissolved	70.6 g CuCl2/100 mL H2O at 0°C
Enthalpy of Solution	Energy change in solution formation	$\Delta H_{solution} = \Delta H_{solute} + \Delta H_{solvent} + \Delta H_{mix}$
Lattice Energy	Energy released in ionic solid formation	NaCl(s) from Na+(g) and Cl-(g)
Hydration Energy	Energy released when ions dissolve	Ca2+(g) + H2O(l) → Ca2+(aq)
Henry's Law	Gas solubility vs. pressure	CO2 in soda

Additional info: Academic context and examples were expanded for clarity and completeness. All images included are directly relevant to the adjacent explanations, visually reinforcing key concepts in solution chemistry.