Solutions and Colloids: Dissolution, Electrolytes, and Solubility

Solutions and Colloids

Review of Solutions

Solutions are homogeneous mixtures composed of two or more substances. The process of forming a solution involves the interaction between solute and solvent molecules, and the resulting mixture can exist in various physical states.

• Solute: The substance being dissolved; it takes on the state of the solvent.

• Solvent: The substance doing the dissolving; typically present in the largest amount.

• Soluble: When a solute dissolves in a solvent (e.g., salt in water).

• Insoluble: When a solute does not dissolve in a solvent (e.g., oil in water).

• Factors affecting solubility:

  1. Nature’s tendency towards mixing (entropy).

  2. Types of intermolecular attractive forces.

Physical States of Solutions:

• Solid solution: Metal alloys (e.g., two or more metals).

• Gas solution: Mixtures of gases (e.g., air).

• Liquid solution: Most commonly discussed in chemistry; water is often the solvent (aqueous solutions).

1. Solutions are homogeneous.

2. The physical state of the solution is typically the same as the solvent.

3. The dissolved solute will not settle out or separate from the solution.

4. Components are dispersed as molecules, atoms, or ions.

Review of Intermolecular Forces (IMFs)

Intermolecular forces are responsible for the interactions between molecules in a solution. Their strength varies among substances and influences solubility.

• IMFs are weaker than covalent bonds.

• Types of IMFs:

  • Dispersion Forces

  • Dipole-Dipole Attractions

  • Hydrogen Bonding

• IMFs are collectively referred to as van der Waals forces.

• Compounds will mix if solvent-solute interactions are equal to or stronger than solvent-solvent and solute-solute interactions.

The Formation of Solutions

Solution formation is governed by three types of intermolecular attractive forces:

• Solute-solute

• Solvent-solvent

• Solute-solvent

The process can be described stepwise:

1. Solute-solute IMFs must be overcome (energy consumed).

2. Solvent-solvent IMFs must be overcome (energy consumed).

3. Solvent-solute attractive forces are established (energy released).

The enthalpy change associated with solution formation is:

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

Solvent-Solvent Interactions

The solubility of substances is often summarized by the rule "Like Dissolves Like". Substances with similar IMFs are likely to be soluble in one another.

• Nonpolar solvents dissolve nonpolar solutes.

• Polar solvents dissolve polar solutes and many ionic solutes.

The Water Molecule

Water is a key solvent in chemistry due to its unique properties:

• Oxygen is more electronegative, pulling electrons toward itself.

• Bent molecular structure with a bond angle of 105°.

• Water is polar, with partial negative charge ($\delta^-$) on O and partial positive charge ($\delta^+$) on H.

• Capable of forming hydrogen bonds, strong electrostatic attractions between H and O.

Electrolytes

Electrolytes are substances that dissolve in water and produce ions, affecting the solution’s conductivity.

• Strong Electrolytes: Nearly 100% of dissolved substance generates ions (e.g., NaCl, HCl).

• Weak Electrolytes: Only a small fraction generates ions (e.g., CH3COOH, MgSO4).

• Nonelectrolytes: Do not produce ions when dissolved (e.g., sugar).

Ionic Electrolytes

Water and other polar molecules are attracted to ions, forming ion-dipole attractions. Water molecules surround and solvate separated ions, a process known as dissociation.

• In ionic lattices, positive and negative ions are held together by electrostatic forces.

• When dissolved, water molecules compete with these forces, pulling ions into solution.

Solubility of Ionic Compounds

The extent of solubility for ionic compounds depends on competing forces:

• Ion-dipole attraction: If dominant, the compound is highly soluble in water.

• Ion-ion attraction: If dominant, the compound remains undissolved and has low solubility.

Solubility

Solubility refers to the maximum concentration of a solute that can be achieved in a particular solvent under given conditions.

• Saturated solution: Solute concentration equals its solubility.

• Unsaturated solution: Solute concentration is less than its solubility.

• Supersaturated solution: Solute concentration is greater than its solubility.

Principles of Solubility

Solubility is limited and depends on several factors:

• Gases are always soluble in each other.

• Liquids that mix in all proportions are miscible (e.g., alcohol and water).

• Liquids that do not mix are immiscible (e.g., oil and water).

• The nature of solvent and solute molecules and their interactions.

• Temperature and pressure (especially for gases).

• Entropy: Nature tends to mix to increase energy dispersal.

Entropy

Entropy is the energy randomization or dispersal in a system. Mixing increases entropy, favoring solution formation.

• In gases, minimal intermolecular interactions allow spontaneous mixing and energy distribution.

Temperature and Gas Solubility in Water

Temperature affects the solubility of gases in water:

• Solubility is given in moles of solute per liter of solution.

• Gases generally have lower solubility than ionic or polar covalent solids.

• For all gases, solubility decreases as temperature increases due to increased kinetic energy.

Pressure and Gas Solubility

Pressure has a major effect on the solubility of gases in liquids. The relationship is described by Henry’s Law:

$C_g = kP_g$

• $C_g$ = concentration of dissolved gas

• $k$ = Henry’s law constant (depends on gas, solvent, and temperature)

• $P_g$ = partial pressure of the gas

Table: Henry's Law Constants for Several Gases in Water at 25°C

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

Deviations from Henry’s Law

Henry’s Law may not apply when a chemical reaction occurs between the gaseous solute and the solvent. For example, ammonia reacts with water to form ammonium and hydroxide ions:

$\ce{NH3 + H2O <=> NH4^+ + OH^-}$

Solutions of Liquids in Liquids

Liquids can dissolve in other liquids, forming miscible or immiscible mixtures:

• Miscible: Mix in all proportions (e.g., ethanol and water).

• Immiscible: Do not mix appreciably (e.g., gasoline and water).

Temperature and Solubility of Solids

Both ionic and covalent solids can dissolve in liquids. Dissolving a solid is usually endothermic, and increased temperature generally increases solubility.

Example equation:

$\ce{NaC2H3O2 (s) ->[H2O] Na^+ (aq) + C2H3O2^- (aq)}$

Supersaturated Solutions

Supersaturated solutions contain more solute than is normally possible at a given temperature. They are prepared by:

• Making a saturated solution at high temperature.

• Slowly cooling to a lower temperature.

• If cooling is slow and careful, the solid will not precipitate, resulting in a supersaturated solution.

Example Problems

• Solvent selection: Choose water or hexane to dissolve benzene, methanol, or sodium amide, and state the intermolecular forces involved.

• Henry’s Law application: Calculate the pressure required to keep CO2 at a certain concentration in club soda.

• Gas solubility calculation: Determine the amount of argon dissolved in water under specified conditions using Henry’s Law.

Additional info: Some context and definitions have been expanded for clarity and completeness.