Review for Exam III: Chemical Bonding and Thermochemistry

Chemical Bonding (Chapter 8)

Lattice Energy and Ionic Compounds

Chemical bonding involves the forces that hold atoms together in compounds. Lattice energy is a key concept in understanding the stability of ionic solids.

• Lattice Energy: The energy required to separate one mole of an ionic solid into its gaseous ions.

• Trends: Lattice energy increases with higher ionic charges and decreases with larger ionic radii.

• Group IA and IIA Metals: Compare their lattice energies to understand periodic trends.

• Born-Haber Cycle: A thermochemical cycle used to calculate lattice energy using enthalpy changes.

Example: NaCl has a higher lattice energy than KCl due to the smaller ionic radius of Na+.

Bonding Models and Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

• Octet Rule: Atoms tend to form bonds until they are surrounded by eight valence electrons.

• Exceptions: Hydrogen (2 electrons), Boron (6 electrons), and expanded octets for elements in period 3 or higher.

• Formal Charge: Used to determine the most stable Lewis structure.

Example: The Lewis structure of CO2 shows double bonds between carbon and each oxygen atom.

Shapes and Polarity (Chapter 9)

VSEPR Theory and Molecular Geometry

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

• Electron Domains: Regions of electron density (bonds and lone pairs) around a central atom.

• Common Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Polarity: Determined by molecular geometry and differences in electronegativity.

Example: H2O is bent and polar due to lone pairs on oxygen.

Bonding Theories (Chapter 9, 9.4-9.5)

Hybridization and Molecular Orbitals

Bonding theories explain how atomic orbitals combine to form molecular orbitals and hybrid orbitals.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (e.g., sp, sp2, sp3).

• Sigma (σ) and Pi (π) Bonds: Sigma bonds are formed by head-on overlap; pi bonds by side-on overlap.

• Bond Order: Number of chemical bonds between a pair of atoms.

Example: Ethylene (C2H4) has a double bond consisting of one sigma and one pi bond.

Thermochemistry (Chapter 5)

Enthalpy and Calorimetry

Thermochemistry studies the heat changes that accompany chemical reactions.

• Enthalpy (ΔH): The heat content of a system at constant pressure.

• Calorimetry: Measurement of heat flow using devices like coffee-cup or bomb calorimeters.

• Specific Heat (c): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Heat Capacity (C): Amount of heat required to raise the temperature of an object by 1°C.

Example: Calculating the heat absorbed by water when a metal is placed in it using $q = mc riangle T$.

Hess's Law and Enthalpy of Formation

Hess's Law states that the total enthalpy change for a reaction is the same, no matter how it is carried out in steps.

• Standard Enthalpy of Formation (ΔHf°): Enthalpy change when one mole of a compound is formed from its elements in their standard states.

• Using Hess's Law: Combine equations to find the enthalpy change for a reaction.

Example: Calculating ΔH for the combustion of methane using enthalpies of formation.

Formulas and Equations

Key Thermochemical Equations

• $q = mc\Delta T$

• $\Delta H = \sum \Delta H_f^{\circ} (\text{products}) - \sum \Delta H_f^{\circ} (\text{reactants})$

• $C = q/\Delta T$

Definitions: Know definitions for bond enthalpy, enthalpy of formation, and related terms.

Additional Info

• Be able to compare bond strengths and predict relative stabilities of compounds.

• Understand the relationship between bond enthalpy and reaction energetics.

• Practice problems involving calculation of enthalpy changes using calorimetry and Hess's Law.

Summary Table: Key Thermochemistry Concepts

Concept	Definition	Formula
Enthalpy (ΔH)	Heat change at constant pressure	$\Delta H = q_p$
Specific Heat (c)	Heat required to raise 1 g by 1°C	$q = mc\Delta T$
Bond Enthalpy	Energy to break one mole of bonds	Sum of bond energies
Standard Enthalpy of Formation (ΔHf°)	Enthalpy change for forming 1 mole from elements	$\Delta H = \sum \Delta H_f^{\circ} (\text{products}) - \sum \Delta H_f^{\circ} (\text{reactants})$