Exam 2 Study Guide: Intermolecular Forces, Solids, and Lattice Energy

Constants and Fundamentals

Important Constants and Conversions

Understanding key constants and conversions is essential for solving problems in general chemistry, especially those involving gases, phase changes, and calculations related to solids.

• Ideal Gas Constant (R): $R = 0.0821\ \text{L} \cdot \text{atm} \cdot \text{mol}^{-1} \cdot \text{K}^{-1} = 8.314\ \text{J} \cdot \text{mol}^{-1} \cdot \text{K}^{-1}$

• Pressure Conversions: $1\ \text{atm} = 760\ \text{mm Hg} = 760\ \text{torr} = 101.3\ \text{kPa} = 1.013\ \text{bar}$

• Avogadro’s Number: $6.022 \times 10^{23}\ \text{mol}^{-1}$

• Temperature Conversion: $K = ^\circ\text{C} + 273.15$

• Density of Water: $1.00\ \text{g/mL}$

Key Equations

• Clausius-Clapeyron Equation: $\ln\left(\frac{P_2}{P_1}\right) = \frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T_1} - \frac{1}{T_2}\right)$

• Bragg’s Law: $n\lambda = 2d\sin\theta$

• Density of a Unit Cell: $\text{Density} = \frac{\text{Mass of atoms in unit cell}}{\text{Volume of unit cell}}$

• Edge Lengths:

  • Body-Centered Cubic (BCC): $a = \frac{4r}{\sqrt{3}}$

  • Face-Centered Cubic (FCC): $a = \sqrt{8}r$

  • Simple Cubic (SC): $a = 2r$

• Coulomb’s Law: $E \propto \frac{q_1 \cdot q_2}{r}$

• Heat Calculations:

  • Boiling/Melting: $q = n\Delta H_{\text{vap/fus}}$

  • Heating: $q = m c \Delta T$

Chapter 11: Intermolecular Forces, Liquids, and Solids

Phases of Matter

The physical properties of matter depend on the phase (state) and the strength of intermolecular forces relative to thermal energy.

State	Density	Shape	Volume	IMF Strength
Gas	Low	Indefinite	Indefinite	Weak
Liquid	High	Indefinite	Definite	Moderate
Solid	High	Definite	Definite	Strong

Shape and Polarity

Molecular shape and polarity are determined using VSEPR theory and AX notation. Polarity affects physical properties and intermolecular forces.

• VSEPR Theory: Predicts molecular geometry based on electron pair repulsion.

• AX Notation: Describes the number of atoms (X) and lone pairs (E) around a central atom (A). Example: AX2E2 for H2O.

• Polarity: Determined by shape and bond dipoles. CO2 is nonpolar (linear, dipoles cancel); H2O is polar (bent, net dipole).

• Polarizability: Larger atoms/molecules are more polarizable, leading to stronger London Dispersion Forces.

Intermolecular Forces (IMFs)

IMFs are forces between molecules that influence physical properties such as boiling point and melting point.

• London Dispersion: Present in all molecules; caused by temporary induced dipoles.

• Dipole-Dipole: Occurs between polar molecules.

• Hydrogen Bonding: Strongest IMF; occurs when H is bonded to F, O, or N ("FON").

• Ion-Dipole: Interaction between an ion and a polar solvent.

Trends: Stronger IMFs lead to higher boiling/melting points, lower vapor pressure, higher viscosity, and greater surface tension.

Properties of Liquids

Liquids exhibit unique properties due to IMFs.

• Surface Tension: Caused by inward forces on surface molecules.

• Adhesion: Attraction between unlike molecules.

• Cohesion: Attraction between like molecules.

• Capillary Action: Movement of liquid in narrow spaces due to adhesion and surface tension.

• Viscosity: Resistance to flow; increases with stronger IMFs.

• Volatility: Tendency to vaporize; decreases with stronger IMFs.

• Dynamic Equilibrium: Occurs when the rate of evaporation equals the rate of condensation.

• Vapor Pressure: Pressure exerted by vapor above a liquid; increases with temperature.

• Boiling Point: Temperature at which vapor pressure equals atmospheric pressure.

Phase Diagrams and Changes

Phase diagrams show the relationship between pressure, temperature, and phases of matter.

• Phase Changes: Sublimation, deposition, fusion (melting), freezing, vaporization, condensation.

• Triple Point: All three phases coexist.

• Critical Point: Marks the supercritical fluid region.

• Phase Boundaries: Lines separating different phases.

• Slope Significance: The slope of the solid-liquid line indicates how pressure affects melting; for water, ice melts under pressure.

Chapter 12: Solids and Modern Materials

Types of Solids

Solids are classified based on their structure and bonding.

• Crystalline Solids: Have long-range order.

• Amorphous Solids: Lack long-range order; random arrangement.

• Classification:

  • Metallic: Sea of electrons; malleable, conductive.

  • Ionic: Held by coulombic forces; hard, brittle.

  • Network Atomic: Covalent bonds throughout (e.g., diamond).

  • Molecular: Held by IMFs; soft, low melting points.

Crystalline Structures

Unit cells are the basic repeating units in crystalline solids.

Type	Atoms/Cell	Coordination Number	Stacking
Simple Cubic (sc)	1	6	aaaa
Body-Centered Cubic (bcc)	2	8	abab
Face-Centered Cubic (fcc)	4	12	abcabc

• Ionic Compounds: Cations and anions form distinct sublattices.

• Tetrahedral and Octahedral Holes: Spaces within the lattice where ions can reside.

Crystal Calculations

Calculations involving unit cells help determine properties like density and packing efficiency.

• Packing Efficiency: Percentage of space occupied by atoms in a unit cell.

• Density Calculations: Relate mass, molar mass, and edge length.

• Bragg’s Law: Used in X-ray diffraction to calculate interatomic distances.

Section 9.11: Lattice Energy

Lattice Energy

Lattice energy is the energy released when gaseous ions form one mole of an ionic solid. It is always negative (exothermic).

• Definition: Energy released upon formation of an ionic solid from gaseous ions.

• Sign: Always negative due to bond formation.

Predicting Lattice Energy

Coulomb’s Law is used to predict lattice energy based on ionic charge and radius.

• Ionic Charge: Higher charges result in higher lattice energy (e.g., CaS vs. KF).

• Ionic Radius: Smaller ions result in higher lattice energy (e.g., LiF vs. CsI).

Example: CaS (+2/-2) has a much higher lattice energy than KF (+1/-1).

Physical Trends

• Melting Point: Higher lattice energy leads to higher melting points and greater hardness/brittleness.

• Solubility: Very high lattice energy can decrease solubility, as ions are strongly attracted and difficult to separate.

Study Strategy

Effective Study Tips

• Practice a wide range of problems to identify weak points.

• Review 3D stacking models to distinguish unit cell arrangements.

• Always check units in calculations, especially temperature (Kelvin) and gas constant (R).