General Chemistry: Molecular Structure, Bonding, Intermolecular Forces, and Solutions

Molecular Structure and Bonding

Lewis Structures and Formal Charge

Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist. Formal charge is a tool used to determine the most stable Lewis structure for a molecule.

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

• Example: In $\mathrm{NH_4^+}$, nitrogen has a formal charge of -1.

Octet Rule and Exceptions

The octet rule states that atoms tend to form compounds in ways that give them eight valence electrons, resembling the electron configuration of a noble gas. Some species violate the octet rule, such as molecules with odd numbers of electrons, electron-deficient molecules, or expanded octets.

• Example: $\mathrm{BF_3}$ violates the octet rule as boron has only six electrons around it.

VSEPR Theory and Molecular Geometry

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

• Common Geometries:

  • Linear: 180° bond angle (e.g., CO2)

  • Trigonal planar: 120° bond angle (e.g., SO3)

  • Bent: <120° or <109.5° (e.g., H2O)

  • Tetrahedral: 109.5° (e.g., CH4)

• Example: SO2 is bent according to VSEPR theory.

Polarity of Molecules

Molecular polarity depends on the difference in electronegativity between atoms and the symmetry of the molecule.

• Nonpolar Molecules: Molecules with symmetrical charge distribution (e.g., CO2).

• Polar Molecules: Molecules with asymmetrical charge distribution (e.g., H2O).

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons.

• sp: Linear geometry (e.g., HCN)

• sp2: Trigonal planar geometry

• sp3: Tetrahedral geometry

Sigma (σ) and Pi (π) Bonds

Sigma bonds are single covalent bonds formed by head-to-head overlap of orbitals. Pi bonds are formed by side-to-side overlap and are present in double and triple bonds.

• Example: In C2H4 (ethylene), there are 5 σ and 1 π bond.

• σ bonds: Allow free rotation; π bonds: Restrict rotation.

Delocalized π Bonding

Delocalized π bonding occurs when π electrons are shared over more than two atoms, as in aromatic compounds.

• Example: C6H6 (benzene) has delocalized π bonding.

Molecular Orbital (MO) Theory

Bond Order

Bond order indicates the strength and stability of a bond. It is calculated as:

• $\text{Bond order} = \frac{1}{2}[(\text{number of bonding electrons}) - (\text{number of antibonding electrons})]$

• Example: O2 has a bond order of 2.

Paramagnetism and Diamagnetism

Paramagnetic molecules have unpaired electrons and are attracted to magnetic fields. Diamagnetic molecules have all electrons paired and are repelled by magnetic fields.

• Example: O2 is paramagnetic.

Bonding and Antibonding Orbitals

• Bonding orbitals: Lower in energy, stabilize the molecule.

• Antibonding orbitals: Higher in energy, destabilize the molecule.

• Electrons in antibonding orbitals can destabilize a molecule.

MO Diagrams

MO diagrams show the relative energy levels of molecular orbitals formed from atomic orbitals. For diatomic molecules like F2, the correct MO diagram shows the filling of σ and π orbitals according to the number of electrons.

Intermolecular Forces

Types of Intermolecular Forces

• London Dispersion Forces: Present in all molecules; arise from temporary dipoles.

• Dipole-Dipole Interactions: Occur between polar molecules.

• Hydrogen Bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

• Ion-Dipole Forces: Occur between ions and polar molecules.

Boiling Point and Intermolecular Forces

• Stronger intermolecular forces lead to higher boiling points.

• Example: NH3 has a higher boiling point than CH4 due to hydrogen bonding.

Surface Tension

Surface tension is the energy required to increase the surface area of a liquid. It is primarily due to hydrogen bonding in water.

Phase Changes and Diagrams

Phase Diagrams

Phase diagrams show the state of a substance at various temperatures and pressures.

• Triple Point: The point where solid, liquid, and gas phases coexist.

• Critical Point: The temperature and pressure above which the liquid and gas phases are indistinguishable.

• Phase Boundaries: Lines separating different phases; the liquid-gas boundary represents vaporization.

Solid-Liquid Line Slope

• For water, the solid-liquid line slopes to the left because ice is less dense than liquid water.

Solutions and Colligative Properties

Concentration Units

• Molality (m): Moles of solute per kilogram of solvent.

Solubility and Saturation

• Unsaturated: Solution can dissolve more solute.

• Saturated: Solution contains the maximum amount of dissolved solute.

• Supersaturated: Solution contains more solute than can normally dissolve at that temperature.

Colligative Properties

Colligative properties depend on the number of solute particles, not their identity.

• Boiling Point Elevation: $\Delta T_b = i K_b m$

• Freezing Point Depression: $\Delta T_f = i K_f m$

• Example: The freezing point of a 1.0 m NaCl solution is calculated using $K_f$ and the van 't Hoff factor $i$ (for NaCl, $i = 2$).

Colloids

Colloids are mixtures where one substance is dispersed evenly throughout another. They are not true solutions and exhibit the Tyndall effect (scattering of light).

• Example: Milk is a colloid.

Summary Table: Types of Intermolecular Forces

Type	Occurs Between	Relative Strength	Example
London Dispersion	All molecules	Weakest	CH4
Dipole-Dipole	Polar molecules	Intermediate	SO2
Hydrogen Bonding	H with N, O, or F	Strongest (of IMFs)	H2O, NH3
Ion-Dipole	Ions and polar molecules	Very strong	Na+ in H2O

Summary Table: Colligative Properties

Property	Definition	Equation
Boiling Point Elevation	Increase in boiling point due to solute	$\Delta T_b = i K_b m$
Freezing Point Depression	Decrease in freezing point due to solute	$\Delta T_f = i K_f m$
Osmotic Pressure	Pressure to stop osmosis	$\Pi = iMRT$

Additional info:

• MO diagrams and their correct filling for F2 and other diatomics are essential for understanding bond order and magnetism.

• Phase diagrams are critical for predicting the behavior of substances under varying temperature and pressure.

• Colligative properties are used to determine molar mass and solution concentration in laboratory settings.