Notes: Liquids and Solids (Chapter 10)

INTERMOLECULAR FORCES (IMFs) AND THE STATES OF MATTER

• The differences in the properties of a solid, liquid, or gas reflect the strengths of the attractive forces between the atoms, molecules, or ions that make up each phase.
• The phase in which a substance exists depends on the relative extents of its intermolecular forces (IMFs) and the kinetic energies (KE) of its molecules.
• If the temperature of a liquid becomes sufficiently low, or the pressure on the liquid becomes sufficiently high, phase transitions can occur.

INTRAMOLECULAR FORCES VS INTERMOLECULAR FORCES

• Intramolecular forces: forces between atoms within a molecule; responsible for bond strength; determine chemical properties.

• Intermolecular forces (IMFs): forces that keep molecules together in a group; generally weaker than intramolecular forces; largely determine physical properties (melting/boiling points, solubility, etc.).

• Visual aid concept: intramolecular vs intermolecular attractions influence whether a substance behaves chemically (reactions) or physically (phase changes, solubility).

• Examples and notes:

  • If a molecule breaks down easily in a reaction, intramolecular forces are weak (bonds within the molecule are weak).
  • Intramolecular forces relate more to chemical properties; IMFs relate more to physical properties.

PHYSICAL VS CHEMICAL PROPERTIES

• Physical properties: can observe or measure without changing chemical composition.
  • Examples: boiling point, melting point, solubility, color, odor, state of matter, density, etc.
• Chemical properties: describe how a substance interacts or changes to form new substances.
  • Examples: toxicity, reactivity, types of chemical bonds formed, oxidation states, flammability, chemical stability, acidity/basicity, radioactivity.
• Key definitions:
  • Physical properties enable us to distinguish substances without changing composition.
  • Chemical properties describe potential chemical changes (conversion to other substances).

INTERMOLECULAR FORCES (IMFs) AND PHYSICAL PROPERTIES

• IMFs are the forces that keep molecules together and influence physical properties such as melting and boiling points.
• Melting point: temperature at which a solid transitions to a liquid.
  • Example: ice melts at approximately T_m = 0^{
    m{o}}C.
• Boiling point: temperature at which a liquid transitions to a gas.
  • Example: water boils at approximately T_b = 100^{
    m{o}}C at 1 atm.
• A compound with a low melting or boiling point indicates weaker IMFs.
• Notes on pressure and temperature references:
  • Normal atmospheric pressure is 1~{
    m atm} = 101~{
    m kPa} (often used in phase diagrams and normal boiling/melting points).

INTERMOLECULAR FORCES TYPES (VAN DER WAALS FORCES)

• All attractive forces between neutral atoms and molecules are collectively known as van der Waals forces (informally called intermolecular attractions).

• Examples of van der Waals forces (in increasing strength):

  • London dispersion forces (dispersion forces)
  • Dipole-dipole interactions
  • Hydrogen bonding

• Polarity and dipoles:

  • Polar covalent bonds lead to dipole moments (partial positive and partial negative charges).
  • Nonpolar molecules have little or no permanent dipole moment, but can experience induced dipoles (and dispersion forces).

DISPERSION FORCES (LONDON FORCES)

• Cause: instantaneous dipoles arising from constant electron motion; these induce dipoles in neighboring molecules.
• Also called London dispersion forces; arise from Coulombic attractions between instantaneous dipoles.
• Characteristics:
  • Present in all molecules (polar and nonpolar).
  • Generally stronger with larger, more polarizable electron clouds and greater molar mass.
  • Often the dominant IMF in nonpolar molecules and noble gases.
• Visual: temporary dipoles lead to weak, short-range attractions; significant only when molecules are very close.

DIPOLE-DIPOLE INTERACTIONS (POLARITY)

• Occur between polar molecules with permanent dipole moments.
• Strength depends on the magnitude of the dipole (larger dipole moments yield stronger attractions).
• Nonpolar vs Polar examples:
  • Nonpolar: noble gases, Br2, I2, N2, Cl2, CO2, CH4, CCl4, etc.
  • Polar: H2O, NH3, SO2, C2H6O (ethanol), etc.

HYDROGEN BONDING

• A special type of dipole-dipole interaction.
• Occurs only when H is bonded to N, O, or F (highly electronegative atoms) in the molecule.
• Rationale: large electronegativity difference between H (~2.1) and the bonded atom (F ≈ 4.0, O ≈ 3.5, N ≈ 3.0) creates highly concentrated partial charges, enabling strong intermolecular attractions.
• Often explains anomalously high boiling points (e.g., water and HF) relative to other compounds in the same group.

ION-DIPOLE INTERACTIONS

• Electrostatic attractions between ions and polar molecules.
• Important for dissolution of ionic compounds in water: water’s polarity stabilizes ions in solution, aiding solvation.

SUMMARY OF IMF TYPES (FLOW DIAGRAM SUMMARY)

• No molecules or ions present? No IMF interactions.
• Are polar molecules involved? If yes, dipole-dipole forces or hydrogen bonding may be present depending on H attached to N, O, or F.
• Are ions present? If yes, ion-dipole interactions come into play.
• Hydrogen bonding occurs if hydrogen is bonded to N, O, or F and polar molecules/ions are involved.
• Examples:
  • London forces only (induced dipoles): Ar, Kr, etc.
  • Dipole-dipole forces: H2S, CH2Cl2, etc.
  • Hydrogen bonding: liquid and solid H2O, NH3, HF.
  • Ion-dipole forces: NaCl, NH4NO3 (ions present with polar solvents).
  • Ionic bonding (section 8.2): e.g., NaCl solid with lattice in solids.

SOLIDS, LATTICES, AND LATTICE ENERGY

• Ionic bonding vs molecular (covalent) bonding:
  • Ionic compounds: no discrete molecules; strong electrostatic forces between metal cations and nonmetal anions; lattice structure instead of individual molecules.
  • Molecular compounds: form discrete molecules; intra- and intermolecular forces differ in strength; IMFs govern physical properties.
• Lattice energy: the energy required to convert a mole of ionic solid to its gaseous ions. For NaCl(s):
  • DH_lattice = $+788 ext{ kJ/mol}$
• Key factors determining lattice energy:
  • Magnitudes of the ionic charges (Q1 and Q2)
  • Distance between ions (d)
  • General trend: lattice energy increases with larger charges and shorter interionic distances.
• Empirical relation (Coulomb’s law intuition):
  • $E<em>{ ext{lattice}} \propto\frac{Q</em>1 Q_2}{d}$
• Example: compare Al2O3 vs Al2Se3 (charges equal, but O^2− is smaller than Se^2−); Al2O3 has larger lattice energy due to shorter distances between ions.

TABLE OF LATTICE ENERGIES AND MELTING POINTS (SELECTED)

• LiF: lattice energy $= 1017 ext{ kJ/mol}$; melting point $T_m ext{ (deg C)} = 845$

• LiCl: lattice energy $= 860 ext{ kJ/mol}$; melting point $T_m = 610$

• LiBr: lattice energy $= 787 ext{ kJ/mol}$; melting point $T_m = 550$

• LiI: lattice energy $= 732 ext{ kJ/mol}$; melting point $T_m = 450$

• NaCl: lattice energy $= 787 ext{ kJ/mol}$; melting point $T_m = 801$

• NaBr: lattice energy $= 736 ext{ kJ/mol}$; melting point $T_m = 750$

• NaI: lattice energy $= 686 ext{ kJ/mol}$; melting point $T_m = 662$

• KCl: lattice energy $= 699 ext{ kJ/mol}$; melting point $T_m = 772$

• KBr: lattice energy $= 689 ext{ kJ/mol}$; melting point $T_m = 735$

• KI: lattice energy $= 632 ext{ kJ/mol}$; melting point $T_m = 680$

• MgCl2: lattice energy $= 2527 ext{ kJ/mol}$; melting point $T_m = 714$

• Na2O: lattice energy $= 2570 ext{ kJ/mol}$; melting point sublimates at high temperature

• MgO: lattice energy $= 3890 ext{ kJ/mol}$; melting point $T_m = 2800$

• ScN: lattice energy $= 7547 ext{ kJ/mol}$; melting point > 3000

• Note: Na2O sublimates at 1275°C

• The magnitude of lattice energy serves as a measure of an ionic compound’s stability; higher lattice energy generally indicates greater stability and higher melting points.

WORKED EXAMPLES (SOLUBILITY, IMF, AND PHASE BEHAVIOR)

Worked Example 1: Solubility predictions based on polarity

• Strategy: assess solute polarity from structure and Lewis structure, apply VSEPR for molecular solutes, predict solubility in polar (water) vs nonpolar (benzene) solvents.

• Predictions:

  • (a) Br2: nonpolar diatomic; more soluble in benzene than in water.
  • (b) NaI: ionic; more soluble in water.
  • (c) Carbon tetrachloride (CCl4): nonpolar; more soluble in benzene.
  • (d) Formaldehyde (CH2O): polar; can form hydrogen bonds in water; more soluble in water.

• Takeaway: polar solutes prefer polar solvents; nonpolar solutes prefer nonpolar solvents; ionic compounds favor water (polar solvents).

Worked Example 2: Order of lattice energies for MgO, CaO, SrO

• Strategy: compare charges (both +2) and interionic distances; lattice energy increases as distance between ions decreases.
• Ionic radii: Mg2+ ≈ 0.72 Å, Ca2+ ≈ 1.00 Å, Sr2+ ≈ 1.18 Å.
• Conclusion:
  • MgO has the smallest distance between ions → largest lattice energy among the three.
  • SrO has the largest distance → smallest lattice energy.
• Result (in increasing lattice energy): SrO < CaO < MgO.
• Provided values: SrO ≈ 3217 kJ/mol, CaO ≈ 3414 kJ/mol, MgO ≈ 3890 kJ/mol.
• Insight: Group 2A metals show ionic radii increase down the group; greater ion separation yields lower lattice energy and typically lower melting points.

PHASE DIAGRAMS

• Phase diagrams show the relationship between phases of a substance as a function of pressure and temperature.
• The diagram combines plots of pressure vs temperature for solid-liquid, liquid-gas, and solid-gas phase equilibria.
• Normal melting point and normal boiling point are read at 1 atm (or 101 kPa).
• Example: Phase diagram of water shows unique features due to hydrogen bonding (e.g., ice is less dense than liquid water, leading to unusual solid-liquid equilibrium behavior at certain pressures).

Phase Diagram of Water (key points)

• Normal melting point occurs at the temperature where solid and liquid phases are in equilibrium at 1 atm.
• Normal boiling point occurs at the temperature where liquid and gas phases are in equilibrium at 1 atm.

Phase Diagram of Carbon Dioxide (CO₂)

• CO₂ can exist as a solid, liquid, or gas depending on T and P.
• Above its critical temperature, CO₂ cannot be liquefied by pressure alone (supercritical fluid, SCF).
• The critical point defines the boundary between gas-like and liquid-like behavior; at sufficiently high T and P, a single supercritical phase exists.
• The critical pressure is the pressure required to liquefy a gas at its critical temperature.

ADDITIONAL TOPICS: WATER, ETHANOL, AND SOLUTIONS

“Like dissolves like” and miscibility

• Substances with similar types and magnitudes of IMFs tend to be soluble in each other (miscible).
• Example: Toluene (C7H8) and Octane (C8H18) are both nonpolar liquids; they are miscible in all proportions.
• Water (polar) and octane (nonpolar) are immiscible due to differing IMF types and magnitudes.
• Ethanol (C2H6O) and water are polar liquids and are miscible.

Miscibility definitions

• Miscible: two liquids that are completely soluble in each other in all proportions.
• Immiscible: two liquids that are insoluble in each other.

Surface Tension

• Surface tension arises from IMFs at the surface pulling molecules toward the bulk; interior pulls are balanced.
• It is the energy required to increase the surface area of a liquid by a unit area: $ext{Surface tension} = \frac{dE}{dA}$.
• Strong IMFs generally yield higher surface tension.
• Water has notably high surface tension due to strong hydrogen bonding.

Capillarity (Capillary Action)

• Cohesion: IMF attraction between like molecules.
• Adhesion: IMF attraction between unlike molecules (e.g., liquid and a solid surface).
• Capillary action occurs when adhesive forces between a liquid and a porous material (or tube) interact with cohesive forces to move liquid upward against gravity.

Viscosity

• Viscosity is a measure of a liquid’s resistance to flow.
• Low viscosity: freely flowing liquids; high viscosity: more resistant to flow.
• Temperature effect: increasing temperature increases KE, helping overcome intermolecular forces, typically reducing viscosity.
• Table 10.2 note: more structurally complex molecules and stronger IMFs lead to higher viscosity.

Vapor Pressure

• The pressure exerted by a vapor in equilibrium with its condensed phase (solid or liquid) at a given temperature.

SUMMARY OF IMF EFFECTS ON PROPERTIES

• Surface tension, capillarity, and viscosity are all influenced by IMF strength.
• Vapor pressure decreases as IMF strength and boiling/melting points increase.
• The greater the polarity and ability to hydrogen bond, the higher the boiling point and the higher the surface tension (e.g., water).

KEY DEFINITIONS AND CONCEPTS (RECAP)

• Intermolecular forces (IMFs): forces that hold molecules together in a substance; examples include London dispersion forces, dipole-dipole interactions, hydrogen bonding, and ion-dipole interactions.
• Intramolecular forces: forces that hold atoms together within a molecule; determine chemical bonding and bond strength.
• Lattice energy: the energy required to convert an ionic solid into gaseous ions; increases with greater ionic charges and shorter interionic distances.
• Phase diagram: map of phase stability as a function of pressure and temperature; normal melting/boiling points are at 1 atm (101 kPa).
• Miscibility: the ability of two liquids to mix in all proportions; related to IMF compatibility ("like dissolves like").
• Critical point: the end point of the phase equilibrium curve between liquid and gas; above this point, the substance exists as a supercritical fluid (SCF).

$E<em>{ ext{lattice}} \propto \frac{Q</em>1 Q_2}{d}$

$\gamma = \frac{dE}{dA}$

• Common critical temperature and pressure examples mentioned: 1 atm = 101 kPa; water and CO₂ phase behavior around these values.

NOTES ABOUT SPECIFIC EXAMPLES (RELEVANT FIGURES IN TEXT)

• Water: high surface tension due to hydrogen bonding; phase diagram shows normal melting/boiling points at 1 atm.
• Carbon dioxide: phase behavior includes a supercritical region above the critical point; SCF exists when T > Tc and P > Pc.
• Sodium chloride vs carbon tetrachloride (NaCl vs CCl4): ionic vs covalent; lattice energy vs melting/boiling properties illustrate IMF roles.
• Lattice energy table (selected values): LiF 1017; LiCl 860; LiBr 787; LiI 732; NaCl 787; NaBr 736; NaI 686; KCl 699; KBr 689; KI 632; MgCl2 2527; Na2O 2570; MgO 3890; ScN 7547 (kJ/mol).