Chapter 14 lecture 1

Solutions: Terminology and Types

Definition and Components of Solutions

A solution is a homogeneous mixture of two or more substances. The solvent is the component present in the greater amount, while the solute is present in the lesser amount. In general chemistry, focus is often placed on aqueous solutions, where water acts as the solvent.

• Solvent: The substance in which the solute dissolves; usually present in larger quantity.

• Solute: The substance dissolved in the solvent; present in smaller quantity.

• Solution: The resulting homogeneous mixture.

Table Purpose: Classification of solutions based on the physical state 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 and Types of Solutions

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature. Solutions are classified as saturated, unsaturated, or supersaturated based on the amount of dissolved solute relative to equilibrium.

• 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; more solute can dissolve.

• Supersaturated Solution: Contains more solute than is normally possible; unstable and can precipitate solute.

Example: Copper(II) chloride in water at 0°C: 70.6 g/100 mL is the solubility. Adding 110 g results in a saturated solution with excess undissolved solute; adding 2 g results in an unsaturated solution.

Energetics of Solution Formation

Why Do Molecular Solutions Form?

The formation of a solution depends on the types of intermolecular forces (IMFs) involved and the entropy (randomness) of the system. Solutions tend to form when solute-solvent attractions are comparable to or greater than solute-solute and solvent-solvent attractions.

• IMFs: The strength and type of interactions (e.g., hydrogen bonding, London dispersion forces) affect solution formation.

• Entropy: Mixing increases disorder, favoring solution formation even if the process is endothermic.

Conditions Favoring Solution Formation

• Weak solute-solute interactions

• Weak solvent-solvent interactions

• Strong solute-solvent interactions

When solute-solvent interactions are stronger than solute-solute or solvent-solvent interactions, solution formation is favored.

Energy Changes During Solution Formation

The process of solution formation involves three steps, each with associated enthalpy changes:

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: Energy released in mixing exceeds energy required to separate solute and solvent. Endothermic: Energy required to separate solute and solvent exceeds energy released in mixing.

Applications of Solution Concepts

• Oil and Water: Immiscible due to weak solute-solvent interactions and strong solvent-solvent interactions.

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

• Heptane and Octane: Both nonpolar hydrocarbons; miscible due to similar LDFs.

• Alcohols and Water: Alcohols (methanol, ethanol, ethylene glycol) are miscible with water due to hydrogen bonding.

Vitamin Solubility: Vitamin C (polar, water-soluble) vs. Vitamin A (nonpolar, fat-soluble).

Trends in Alcohol Solubility

As the hydrocarbon chain in alcohols increases, solubility in water decreases due to reduced polarity.

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 (ion-dipole attraction). Each ion is surrounded by water molecules, which prevents recombination.

Enthalpy of Hydration and Lattice Energy

The enthalpy of solution for ionic compounds is:

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

Lattice energy ($\Delta H_{lattice}$) is the energy released when ions form an ionic solid from gaseous ions. Enthalpy of hydration ($\Delta H_{hydration}$) is the energy released when ions are hydrated in water.

• Higher charge density (charge/volume) leads to more exothermic hydration.

• Higher lattice energy makes a compound less soluble in water.

Summary of Lattice Energy

• Greater ionic charge increases lattice energy.

• Lattice energy decreases with increasing ionic radius.

• Higher lattice energy reduces solubility in water.

Effect of Temperature and Pressure on Solubility

Solubility of most solid solutes increases with temperature, but some exceptions exist (e.g., Na2SO4).

Solubility of gases decreases with increasing temperature (thermal pollution). Solubility of gases increases with pressure, described by Henry's Law:

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

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

Example: Solubility of O2 in water at 25°C, $P_{O_2} = 0.21$ atm, $k_H = 1.28 \times 10^{-3}$ mol L-1 atm-1:

$S_{O_2} = k_H P_{O_2} = 1.28 \times 10^{-3} \times 0.21 = 2.7 \times 10^{-4}$ mol/L

Key Concepts and Summary

• Solutions are classified by the physical state of solute and solvent.

• Solubility depends on intermolecular forces, entropy, and energetics.

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

• Ionic solutions involve lattice energy and enthalpy of hydration.

• Temperature and pressure affect solubility of solids and gases differently.