Intermolecular Forces and Solution Formation: Chapter 13 Study Notes

Intermolecular Forces

Introduction to Intermolecular Forces

Intermolecular forces are the attractive or repulsive forces that exist between molecules, influencing physical properties such as boiling point, melting point, solubility, and viscosity. These forces explain why some solids dissolve in water, why noble gases can liquefy, and why substances exist as gases, liquids, or solids at room temperature.

• Intermolecular forces may be attractive or repulsive.

• Johannes D. van der Waals first postulated intermolecular forces to explain the properties of real gases.

Types of Intermolecular Forces

• van der Waals forces include:

  • London (dispersion) forces

  • Dipole-dipole forces

  • Dipole-induced dipole forces

• Other intermolecular forces:

  • Ion-dipole interactions

  • Ion-induced dipole interactions

  • Hydrogen bonding

London (Dispersion) Forces

Origin and Properties

London forces arise from temporary variations in electron density in atoms and molecules. At any instant, the electron distribution may be unsymmetrical, producing an instantaneous dipole that can induce a dipole in neighboring molecules.

• Present between all molecules, whether polar or nonpolar.

• Stronger in molecules that are easily polarizable.

• Larger and heavier molecules exhibit stronger dispersion forces than smaller and lighter ones.

Example: Neon atoms do not have a permanent dipole, but can exhibit dispersion forces.

Magnitude of Dispersion Forces

• Cylindrical-shaped molecules (e.g., n-pentane) have greater surface area and stronger dispersion forces than spherical-shaped molecules (e.g., neopentane).

Dipole-Dipole and Dipole-Induced Dipole Forces

Dipole-Dipole Forces

These forces occur between molecules with permanent dipoles. The positive end of one molecule is attracted to the negative end of another.

Dipole-Induced Dipole Forces

These operate between polar molecules (with permanent dipoles) and nonpolar molecules. The polar molecule induces a dipole in the neighboring nonpolar molecule.

• Interaction energy depends on the dipole moment of the polar molecule.

Examples:

• Oxygen gas dissolves in water because water's permanent dipole induces a dipole in oxygen.

• Nonpolar iodine dissolves in polar ethanol due to dipole-induced dipole forces.

Ion-Dipole and Ion-Induced Dipole Interactions

Ion-Dipole Interactions

These occur between ions and polar molecules. The strength depends on the charge and size of the ion and the magnitude of the dipole moment of the polar molecule.

• Hydration of ions: Sodium ions (Na+) and chloride ions (Cl-) are surrounded by water molecules in solution, stabilizing them through ion-dipole interactions.

Hydrogen Bonding

Definition and Conditions

Hydrogen bonding is an electrostatic force of attraction between a covalently bonded hydrogen atom (attached to a highly electronegative atom such as fluorine, oxygen, or nitrogen) and the electronegative atom of another molecule.

• A hydrogen atom attached to a relatively electronegative atom is a hydrogen bond donor.

• The electronegative atom (F, O, N) is a hydrogen bond acceptor.

Effects of Hydrogen Bonding

• Higher boiling point in ethanol compared to diethyl ether.

• Higher viscosity of sulfuric acid and glycerol.

• Water is a liquid at room temperature, whereas H2S is a gas.

• Formation of dimers in carboxylic acid molecules.

• Cage-like structure of water molecules in ice causes lower density of ice.

Comparison of Intermolecular Forces

Relative Strengths and Examples

Type	Model	Interaction	Energy (kJ/mol)	Example
Ion-dipole	Ion + dipole	Charge-dipole	40–600	Na+ in H2O
H bond	Dipole + dipole	Polar H bond	10–40	H2O, HF
Dipole-dipole	Dipole + dipole	Dipole-dipole	3–20	HCl, CH3Cl
Ion-induced dipole	Ion + induced dipole	Charge-induced dipole	3–15	Fe2+ in O2
Dispersion (London)	Induced dipole + induced dipole	Polarizable	1–10	H2, I2

Solutions: Effect of Intermolecular Forces

Mixing Solute and Solvent

For a solution to form, solute-solute and solvent-solvent attractive forces must be overcome (endothermic), and new solute-solvent attractions must be formed (exothermic).

• Both processes are endothermic, but the formation of new attractions is exothermic.

• Oil and water do not mix because oil is nonpolar and water is polar. "Like dissolves like" principle applies.

Common Types of Solutions

Solute Phase	Solvent Phase	Example
Gas	Gas	Air (O2 and N2)
Gas	Liquid	Carbonated water (CO2 in H2O)
Liquid	Liquid	Vodka (ethanol and water)
Solid	Liquid	Seawater (NaCl in water)
Solid	Solid	Brass (Cu and Zn)

Solution Process and Enthalpy of Solution

Steps in Solution Formation

1. Separating the solute into its constituent particles ($\Delta H_{solute}$, endothermic)

2. Separating the solvent particles from each other ($\Delta H_{solvent}$, endothermic)

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

The overall enthalpy change ($\Delta H_{solution}$) is:

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

• If the sum of endothermic terms is approximately equal to the exothermic term, $\Delta H_{solution}$ is about zero.

• If the sum of endothermic terms is less than the exothermic term, $\Delta H_{solution}$ is negative (exothermic).

• If the sum of endothermic terms is greater than the exothermic term, $\Delta H_{solution}$ is positive (endothermic).

Solubility Limit and Saturation

• Saturated solution: Contains the maximum amount of solute that can dissolve at a given temperature.

• Unsaturated solution: Contains less solute than saturation.

• Supersaturated solution: Contains more solute than saturation; unstable.

Temperature and Pressure Dependence of Solubility

Solids in Water

• Solubility of most solids increases with temperature.

• Solubility curves can predict saturation, unsaturation, or supersaturation.

Gases in Water

• Solubility of gases decreases with increasing temperature.

• Solubility of gases increases with increasing pressure.

Henry's Law: The solubility of a gas ($S_{gas}$) is directly proportional to its partial pressure ($P_{gas}$) at a given temperature.

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

Where $k_H$ is Henry's law constant.

Solution Concentration Terms

Unit	Definition
Molarity (M)	Amount solute (in mol) / volume solution (in L)
Molality (m)	Amount solute (in mol) / mass solvent (in kg)
Mole fraction (x)	Amount solute (in mol) / total amount of solute and solvent (in mol)
Mass percent (%)	Mass solute / mass solution × 100%
Parts per million (ppm)	Mass solute / mass solution × 106
Parts per billion (ppb)	Mass solute / mass solution × 109

Summary Table: Common Laboratory Solvents

Common Polar Solvents	Common Nonpolar Solvents
Water (H2O)	Hexane (C6H14)
Acetone (CH3COCH3)	Diethyl ether (CH3CH2OCH2CH3)
Methanol (CH3OH)	Toluene (C6H5CH3)
Ethanol (CH3CH2OH)	Carbon tetrachloride (CCl4)

Key Equations

• Enthalpy of Solution: $\Delta H_{solution} = \Delta H_{solute} + \Delta H_{solvent} + \Delta H_{mix}$

• Henry's Law: $S_{gas} = k_H P_{gas}$

Examples and Applications

• Oil and water: Do not mix due to polarity differences.

• Solubility curves: Used to determine saturation status at different temperatures.

• Hydrogen bonding: Explains water's high boiling point and ice's low density.

Additional info: These notes expand on the original slides and images by providing definitions, explanations, and tables for clarity and completeness, suitable for exam preparation in General Chemistry.