Intermolecular Forces, Types of Solids, and Properties of Gases: Study Notes

Intermolecular Forces (IMFs)

Definition and Importance

Intermolecular forces (IMFs) are the attractions between molecules. IMFs explain why substances exhibit specific properties such as melting and boiling points, or exist in different states of matter at the same temperature. IMFs are weaker than chemical bonds.

Types of Intermolecular Forces

• London Dispersion Forces (LDFs):

  • Present in all molecules; all molecules can subject each other to these forces.

  • LDFs are the only force of attraction in nonpolar substances (e.g., H2, N2, He, Ne, CH4, C6H6, etc.).

  • Electron clouds can shift around and form a temporary dipole. This temporary dipole can induce another molecule’s electron cloud to form an induced dipole.

  • Electrostatic attraction will form between these temporary dipoles.

  • LDFs are a relatively weak force.

  • The larger the electron cloud, the more polarizable a molecule will be. The ability to form an induced dipole is known as polarizability.

  • LDFs are stronger in molecules with:

    • Larger electron clouds (more polarizable)

    • Greater surface area (long, straight molecules have stronger LDFs than compact, branched ones)

  • LDFs can be enhanced by π bonding (double or triple bonds).

• Dipole-Dipole Forces:

  • Occur in molecules that have a permanent dipole (polar molecules).

  • Stronger than LDFs, as these dipoles are permanent.

  • Strength increases as the overall dipole moment increases.

  • A greater difference in electronegativity between atoms in a bond means a greater dipole moment.

  • Molecules with dipole-dipole attractions also have LDFs.

• Dipole-Induced Dipole Forces:

  • Occur between a polar and a nonpolar molecule.

  • Strength depends on the magnitude of the dipole and the polarizability of the nonpolar molecule.

• Ion-Dipole Forces:

  • Occur between ions and polar molecules.

  • Stronger than hydrogen bonding.

• Hydrogen Bonding:

  • A type of dipole-dipole attraction, but stronger.

  • Not an actual bond, but a strong attraction between molecules.

  • Occurs in molecules that contain H–F, H–O, or H–N bonds.

  • Strength order: H–F > H–O > H–N.

  • Water (H2O) has especially strong hydrogen bonding due to its ability to form multiple hydrogen bonds.

Types of Solids

Classification of Solids

Solids are classified by their lattice structure and the strength of the forces holding them together. There are four main types:

• Molecular Solids:

  • Weakest type of solid; held together by weak IMFs.

  • Made up of individual units of covalently bonded molecules.

  • Not conductive; usually large molecules or polymers (e.g., glucose).

  • Solid water (ice) is a molecular solid held together by hydrogen bonding.

  • Generally have low melting points and high vapor pressures.

• Ionic Solids:

  • Crystal lattices made up of positively and negatively charged ions.

  • Strong electrostatic attractions; high melting and boiling points.

  • Brittle; conduct electricity when ions are mobile (molten or dissolved).

  • Generally between an ionically bonded metal and a nonmetal.

• Metallic Solids:

  • Lattice of positively charged metal cations and a "sea" of delocalized valence electrons.

  • High melting and boiling points, low vapor pressure.

  • Good conductors of heat and electricity; malleable and ductile.

  • Can form alloys with other metals:

    • Interstitial alloys: Small atoms (e.g., H, B, C, N) fit between metal atoms.

    • Substitutional alloys: Metal atoms of similar radii replace each other in the lattice.

• Covalent Network Solids:

  • Atoms are covalently bonded in a continuous network (e.g., diamond, graphite, SiO2, BN, SiC).

  • Very hard, strong, and have high melting points.

  • Rigid and brittle due to fixed covalent bond angles.

  • Generally only formed from nonmetals.

Properties of Solids, Liquids, and Gases

Particulate Diagrams and States of Matter

• Solids: Definite shape and volume; particles vibrate in place; not compressible.

• Liquids: Definite volume, indefinite shape; particles move more freely; not compressible.

• Gases: Indefinite shape and volume; particles move freely and are very compressible.

Crystalline vs. Amorphous Solids

• Crystalline solids: Have ordered 3D lattices.

• Amorphous solids: Have disordered 3D lattices.

Properties of Liquids

• Liquids cannot be compressed due to the lack of space between molecules.

• Properties determined by IMFs:

  • Boiling and melting point

  • Viscosity: Stronger IMFs = higher viscosity (resistance to flow).

  • Surface tension: Stronger IMFs = higher surface tension.

  • Capillary action: Stronger IMFs = increased capillary action.

Properties of Gases

• Gases have gained enough energy to completely break free of their IMFs and move freely.

• Very compressible due to the distances between particles.

• Shape and volume are dependent on temperature and pressure.

• Frequency and strength of collisions are dependent on temperature, pressure, and volume.

Ideal Gas Laws

Relationships Between Gas Properties

• Boyle’s Law: At constant temperature, pressure and volume are inversely proportional.

  • $P_1V_1 = P_2V_2$

• Charles’s Law: At constant pressure, volume and temperature are directly proportional.

  • $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

• Gay-Lussac’s Law: At constant volume, pressure and temperature are directly proportional.

  • $\frac{P_1}{T_1} = \frac{P_2}{T_2}$

• Avogadro’s Law: At constant temperature and pressure, volume and moles are directly proportional.

  • $\frac{V_1}{n_1} = \frac{V_2}{n_2}$

Ideal Gas Law

• Combines all the above relationships:

$PV = nRT$

• P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K)

• R values:

  • 0.08206 (L·atm)/(mol·K)

  • 62.36 (L·torr)/(mol·K)

  • 8.314 (L·kPa)/(mol·K)

Gas Density and Molar Mass

• Density: $d = \frac{MP}{RT}$

• Molar mass: $M = \frac{dRT}{P}$

Dalton’s Law of Partial Pressures

• The total pressure of a mixture of gases equals the sum of the partial pressures of each gas.

• $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

• When gases are collected "over water," subtract the vapor pressure of water from the total pressure to get the pressure of the dry gas.

Summary Table: Types of Solids

Type of Solid	Constituents	Forces Holding Together	Properties	Examples
Molecular	Molecules	IMFs (LDF, dipole-dipole, H-bonding)	Low melting point, not conductive	Ice, glucose
Ionic	Cations and anions	Electrostatic (ionic) bonds	High melting point, brittle, conductive when molten/dissolved	NaCl, KBr
Metallic	Metal atoms	Metallic bonding (sea of electrons)	High melting point, conductive, malleable	Fe, Cu, alloys
Covalent Network	Nonmetal atoms	Covalent bonds	Very high melting point, hard, not conductive	Diamond, SiO2

Additional info: These notes expand on the original content by providing definitions, examples, and formulas for clarity and completeness. The summary table is inferred from the context of the notes and standard chemistry curricula.