General Chemistry II: Practice Exam 2 Study Guide

Intermolecular Forces and Lattice Energy

Types of Intermolecular Forces

Intermolecular forces are the forces of attraction between molecules, which influence physical properties such as boiling and melting points.

• Dispersion Forces (London Forces): Weak forces present in all molecules, especially significant in nonpolar molecules due to temporary dipoles.

• Dipole-Dipole Interactions: Occur between polar molecules with permanent dipoles.

• Hydrogen Bonding: A strong type of dipole-dipole interaction occurring when hydrogen is bonded to highly electronegative atoms (N, O, F).

• Ionic Bonds: Electrostatic attractions between oppositely charged ions, much stronger than typical intermolecular forces.

Example: CH4 has a lower boiling point than SiH4, GeH4, and SnH4 because it has the weakest London dispersion forces due to its smaller size and lower polarizability.

Lattice Energy

Lattice energy is the energy required to separate one mole of an ionic solid into its gaseous ions. It is a measure of the strength of the ionic bonds in a crystalline solid.

• Factors Affecting Lattice Energy:

  • Charge of ions: Higher charges increase lattice energy.

  • Size of ions: Smaller ions increase lattice energy due to closer packing.

Example: NaF has a higher lattice energy than NaCl, CaI2, or CsBr due to the small size and high charge of Na+ and F-.

Isoelectronic Species and Periodic Trends

Isoelectronic Series

Isoelectronic species have the same number of electrons but different nuclear charges. Their relative sizes depend on the number of protons: more protons pull electrons closer, resulting in a smaller radius.

• Order of Size: For isoelectronic species, the one with the most protons is the smallest.

Periodic Trends

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons. Increases across a period.

• Atomic Radius: Decreases across a period (left to right), increases down a group.

• Ionization Energy: Energy required to remove an electron. Increases across a period, decreases down a group.

• Electron Affinity: Energy change when an electron is added to a neutral atom. Becomes more negative across a period.

Example: Bromine has a higher effective nuclear charge and a smaller atomic radius than selenium.

Bonding and Lewis Structures

Bond Strength and Order

• Bond Order: The number of chemical bonds between a pair of atoms. Higher bond order means stronger and shorter bonds.

• Example: A triple bond (C≡C) is stronger than a double bond (C=C) or a single bond (C–C).

Formal Charge

Formal charge is used to determine the most likely Lewis structure for a molecule.

• Formula: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Example: In SO22-, the formal charge on sulfur can be calculated using the above formula.

Lewis Structures

• Show all valence electrons as dots or lines (bonds).

• Minimize formal charges for the most stable structure.

Electron Configuration

Condensed Electron Configuration

Electron configuration describes the distribution of electrons in atomic orbitals.

• Condensed Notation: Uses the previous noble gas in brackets, followed by the remaining configuration.

• Example: Phosphorus (P): [Ne] 3s2 3p3

Molecular Geometry and Hybridization

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shape of molecules based on electron pair repulsion.

• Electron-Domain Geometry: Arrangement of electron domains (bonding and lone pairs) around the central atom.

• Molecular Geometry: Arrangement of atoms (ignoring lone pairs).

Hybridization

• sp3: Tetrahedral geometry, 4 electron domains.

• sp2: Trigonal planar geometry, 3 electron domains.

• sp: Linear geometry, 2 electron domains.

Example: In PF3, phosphorus is sp3 hybridized with a trigonal pyramidal molecular geometry.

Bond Angles

• Tetrahedral: 109.5°

• Trigonal Planar: 120°

• Linear: 180°

Polarity

Molecular polarity depends on the difference in electronegativity and the geometry of the molecule.

• Polar Molecules: Have a net dipole moment due to asymmetrical charge distribution.

• Nonpolar Molecules: Symmetrical molecules with no net dipole moment.

Reference Tables

Common Constants and Equations

• Speed of light: $c = 2.998 \times 10^8 \text{ m/s}$

• Planck's constant: $h = 6.626 \times 10^{-34} \text{ J·s}$

• Energy of a photon: $E = h\nu$

• Relationship between wavelength and frequency: $c = \lambda \nu$

• Rydberg equation for hydrogen: $\frac{1}{\lambda} = R_H \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$

Electronegativity Table

Electronegativity values are used to predict bond polarity and molecular polarity.

Element	Electronegativity
H	2.1
F	4.0
O	3.5
N	3.0
Cl	3.0
Br	2.8
I	2.5

Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

• Groups: Vertical columns with similar valence electron configurations.

• Periods: Horizontal rows indicating energy levels.

Summary Table: Molecular Geometry and Hybridization

Example Molecule	Electron-Domain Geometry	Molecular Geometry	Hybridization	Bond Angle
PF3	Tetrahedral	Trigonal Pyramidal	sp3	~107°
TeBr4	Trigonal Bipyramidal	See-Saw	sp3d	~90°, 120°, 180°