Chapter 13: Solutions – Structure, Properties, and Colligative Effects

Solutions: Structure and Properties

Introduction to Solutions

Solutions are homogeneous mixtures composed of two or more substances. The process of solution formation is governed by the interactions between solute and solvent particles, and the tendency of nature toward spontaneous mixing (entropy).

• Solute: The substance being dissolved.

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

• Homogeneous mixture: Uniform composition throughout.

• Spontaneous mixing: Driven by entropy, leading to uniform distribution of solute.

Example: Wood frogs survive freezing temperatures by flooding their blood with glucose, lowering the freezing point of their body fluids.

Common Types of Solutions

Solutions can exist in various combinations of solute and solvent phases. The table below summarizes common types:

Solution Phase	Solute Phase	Solvent Phase	Example
Gaseous solution	Gas	Gas	Air (O2 and N2)
Liquid solution	Gas	Liquid	Club soda (CO2 in water)
Liquid solution	Liquid	Liquid	Vodka (ethanol and water)
Solid solution	Solid	Solid	Brass (copper and zinc)

Solubility and Miscibility

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent. Miscible liquids are mutually soluble, while immiscible liquids are not.

• Soluble: Salt in water.

• Insoluble: Oil in water.

• Miscible: Alcohol and water.

• Immiscible: Oil and water.

Nature’s Tendency Toward Mixing: Entropy

Entropy and Solution Formation

Mixing increases entropy, which is the measure of energy dispersal in a system. Even when potential energy does not decrease, mixing occurs spontaneously due to entropy.

• Entropy: Drives spontaneous mixing.

• Example: Mixing of neon and argon gases after removing a barrier.

Intermolecular Forces and Solution Formation

Types of Intermolecular Forces

The formation of solutions depends on the types and strengths of intermolecular forces between solute and solvent particles.

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

• Dipole-dipole forces: Between polar molecules.

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

• Ion-dipole forces: Between ions and polar molecules.

Solution Interactions

To form a solution, solute–solute and solvent–solvent interactions must be overcome, and new solute–solvent interactions must be formed.

• Solute–solute interactions: Must be broken (endothermic).

• Solvent–solvent interactions: Must be broken (endothermic).

• Solute–solvent interactions: Formed (exothermic).

Relative Interactions and Solution Formation

The table below summarizes how the relative strengths of interactions affect solution formation:

Solvent–solute interactions	Solvent–solvent and solute–solute interactions	Result
>		Solution generally forms
=		Solution generally forms
<		Solution may or may not form

Solubility of Vitamins: Polar vs. Nonpolar

Classifying Vitamins by Solubility

Vitamins can be classified as water-soluble or fat-soluble based on their molecular structure and polarity.

• Water-soluble: Highly polar, can hydrogen-bond with water (e.g., Vitamin C, Vitamin B5).

• Fat-soluble: Mostly nonpolar, dissolve in fats (e.g., Vitamin K3, Vitamin A).

Energetics of Solution Formation

Enthalpy of Solution

The overall enthalpy change for making a solution depends on three steps:

• 1. ΔHsolute: Overcome attractions between solute particles (endothermic).

• 2. ΔHsolvent: Overcome attractions between solvent particles (endothermic).

• 3. ΔHmix: Form new attractions between solute and solvent (exothermic).

Overall equation:

$\Delta H_{\text{soln}} = \Delta H_{\text{solute}} + \Delta H_{\text{solvent}} + \Delta H_{\text{mix}}$

Exothermic vs. Endothermic Solution Formation

If the energy released in forming solute–solvent attractions is greater than the energy required to break solute–solute and solvent–solvent attractions, the process is exothermic. If less, it is endothermic.

Heats of Hydration

For ionic compounds in water, the heat of hydration combines the energy to overcome water–water attractions and the energy released in forming ion–water attractions.

• Lattice energy: Energy holding ions together in a solid.

• Heat of hydration: Energy released when ions dissolve in water.

Solution Equilibrium and Solubility Limit

Dynamic Equilibrium

When a solution reaches saturation, the rate of dissolution equals the rate of deposition, resulting in dynamic equilibrium.

• Saturated: Maximum solute dissolved.

• Unsaturated: Less than maximum solute.

• Supersaturated: More than maximum solute; unstable.

Purification by Recrystallization

Recrystallization is a method to purify solids by dissolving them in hot solvent and allowing them to crystallize upon cooling, leaving impurities behind.

Temperature and Pressure Dependence of Solubility

Solubility of Solids and Gases

Solubility of solids generally increases with temperature, while solubility of gases decreases with temperature.

Supersaturation and Soda Pop Example

Supersaturated solutions can be formed by changing temperature or pressure. For example, warm soda pop fizzes more because CO2 is less soluble at higher temperatures.

Pressure Dependence: Henry’s Law

The solubility of a gas in a liquid is directly proportional to its partial pressure above the liquid, described by Henry’s Law:

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

Gas	kH (M/atm)
O2	1.3 × 10–3
N2	6.1 × 10–4
CO2	3.4 × 10–2
NH3	5.8 × 101
He	3.7 × 10–4

Solution Concentration Units

Common Concentration Units

Concentration describes the amount of solute in a given amount of solution. Common units include molarity, molality, mass percent, volume percent, ppm, and ppb.

Unit	Definition	Units
Molarity (M)	amount solute (mol) / volume solution (L)	mol/L
Molality (m)	amount solute (mol) / mass solvent (kg)	mol/kg
Mole fraction (χ)	moles of component / total moles	None
Mass percent (%)	mass solute / mass solution × 100%	%
Parts per million (ppm)	mass solute / mass solution × 106	ppm
Parts per billion (ppb)	mass solute / mass solution × 109	ppb

Preparing Solutions

To prepare a solution of known concentration, calculate the required mass of solute and dissolve it in solvent to reach the desired volume.

Molarity and Dissociation

Molarity allows calculation of the concentration of ions in solution after dissociation. For example, 1.0 M CaCl2 yields 1.0 M Ca2+ and 2.0 M Cl–.

Percent, ppm, and ppb Calculations

Percent, ppm, and ppb are used for very dilute solutions. The formulas are:

$\text{Percent} = \frac{\text{Part (solute)}}{\text{Whole (solution)}} \times 100\%$

$\text{PPM} = \frac{\text{Part (solute)}}{\text{Whole (solution)}} \times 10^6$

$\text{PPB} = \frac{\text{Part (solute)}}{\text{Whole (solution)}} \times 10^9$

Colligative Properties

Definition and Importance

Colligative properties depend only on the number of solute particles, not their identity. These include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.

Vapor Pressure Lowering and Raoult’s Law

The vapor pressure of a solvent in a solution is always lower than that of the pure solvent. Raoult’s Law describes this relationship:

$P_{\text{solvent in solution}} = \chi_{\text{solvent}} \cdot P^{\circ}_{\text{solvent}}$

Freezing Point Depression

The freezing point of a solution is lower than that of the pure solvent. The change is proportional to the molality:

$\Delta T_f = m \cdot k_f$

Boiling Point Elevation

The boiling point of a solution is higher than that of the pure solvent. The change is proportional to the molality:

$\Delta T_b = m \cdot k_b$

Osmosis and Osmotic Pressure

Osmosis is the flow of solvent from low to high concentration through a semipermeable membrane. Osmotic pressure is given by:

$\Pi = M R T$

Van’t Hoff Factor

The van’t Hoff factor, i, accounts for the number of particles produced by dissociation of ionic compounds:

$i = \frac{\text{moles of particles in solution}}{\text{moles of formula units dissolved}}$

Colligative Properties in Medical Solutions

Isosmotic, hyperosmotic, and hyposmotic solutions affect cells by controlling the net flow of water, which is critical in medical applications.

Additional info: This study guide covers all major aspects of Chapter 13: Solutions, including structure, properties, energetics, concentration units, and colligative effects, with relevant examples and formulas for exam preparation.