General Chemistry Exam Review: Atomic Structure, Bonding, Molecular Geometry, and Chemical Properties

Atomic Structure and Periodic Trends

Lattice Energy and Ionic Compounds

Lattice energy is the energy released when ions in the gas phase combine to form an ionic solid. It depends on the charges of the ions and the distance between them (ionic radii).

• Key Formula: $|LE| \propto \dfrac{q_1 q_2}{r}$, where $q_1$ and $q_2$ are the charges and $r$ is the distance between ions.

• Trend: Higher charges and smaller ionic radii result in larger lattice energies.

• Example: MgO has a higher lattice energy than NaBr due to higher charges and smaller radii.

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in the gas phase. The second ionization energy is always higher than the first, especially when removing an electron from a noble gas configuration.

• Key Point: Atoms with a stable noble gas configuration after the first ionization have very high second ionization energies.

• Example: Na has a very high second ionization energy because after losing one electron, it achieves a noble gas configuration.

Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond.

• Trend: Increases across a period (left to right), decreases down a group.

• Order Example: F > P > Mg > Cs

Chemical Bonding and Molecular Structure

Resonance Structures

Resonance structures are different Lewis structures for the same molecule, showing delocalization of electrons.

• Key Point: The number of resonance structures depends on the possible locations of double bonds and lone pairs.

• Example: XeO3 can have multiple resonance structures with formal charges distributed among the oxygen atoms.

Lewis Structures and Formal Charge

Lewis structures represent the arrangement of electrons in a molecule. Formal charge helps determine the most stable structure.

• Formula: $\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \dfrac{1}{2} \times \text{Bonding electrons}$

• Application: Structures with minimized formal charges are generally more stable.

Molecular Orbital Theory

Molecular orbital (MO) theory describes the combination of atomic orbitals to form molecular orbitals, which can be bonding or antibonding.

• Bonding MO: Constructive overlap of atomic orbitals increases electron density between nuclei.

• Antibonding MO: Destructive overlap creates a node between nuclei.

• Example: $\sigma_{pp}$ is formed by head-to-head overlap of p orbitals.

Hybridization and Molecular Geometry

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

• sp: Linear geometry

• sp2: Trigonal planar geometry

• sp3: Tetrahedral geometry

• Example: The carbon in a carboxylic acid group is sp2 hybridized and trigonal planar.

Chemical Reactions and Stoichiometry

Combustion Analysis

Combustion analysis is used to determine the empirical formula of hydrocarbons by measuring the amounts of CO2 and H2O produced.

• Key Steps: Calculate moles of C and H from CO2 and H2O, then determine the percentage of pure hydrocarbon.

Bond Energy

Bond energy is the energy required to break a bond in a molecule.

• Example: The bond energy of Br–Br is 193 kJ/mol, corresponding to the reaction $Br_2(g) \rightarrow 2Br(g)$.

Stoichiometry and Limiting Reactant

Stoichiometry involves calculating the amounts of reactants and products in chemical reactions.

• Example: Calculating the mass of chlorine in a sample of C2F4Cl2 using molar mass and percent composition.

Molecular Polarity and Geometry

Polarity of Molecules

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

• Nonpolar: Symmetrical molecules with equal bond dipoles (e.g., CO2).

• Polar: Asymmetrical molecules or those with lone pairs on the central atom (e.g., SF4).

Electron Group Geometry

The geometry around a central atom is determined by the number of electron groups (bonding and lone pairs).

• Linear: 2 electron groups

• Trigonal planar: 3 electron groups

• Tetrahedral: 4 electron groups

• Trigonal bipyramidal: 5 electron groups

• Octahedral: 6 electron groups

Types of Chemical Bonds

Ionic vs. Covalent Bonds

Ionic bonds form between metals and nonmetals with large differences in electronegativity, while covalent bonds form between nonmetals with similar electronegativity.

• Percent Ionic Character: Greatest when the difference in electronegativity is largest (e.g., H–F bond).

Sigma and Pi Bonds

Sigma (σ) bonds are formed by head-to-head overlap of orbitals, while pi (π) bonds are formed by side-to-side overlap.

• Single bond: 1 σ bond

• Double bond: 1 σ and 1 π bond

• Triple bond: 1 σ and 2 π bonds

• Example: Caffeine contains 21 σ bonds and 4 π bonds.

Organic Chemistry Basics

Functional Groups Containing Nitrogen

Organic families that necessarily contain nitrogen include amines and amides.

• Amines: R–NH2, R2NH, R3N

• Amides: R–CO–NH2

Tables

Hybridization and Molecular Geometry Table

This table summarizes the relationship between hybridization and molecular geometry for central atoms in organic functional groups.

Hybridization	Molecular Geometry
sp	Linear
sp2	Trigonal planar
sp3	Tetrahedral
sp3d	Trigonal bipyramidal
sp3d2	Octahedral

Sigma and Pi Bonds in Caffeine

The following table shows the number of sigma and pi bonds in caffeine:

Type of Bond	Number
σ-bonds	21
π-bonds	4

Additional info:

• Some explanations and context have been expanded for clarity and completeness.

• All major topics from the provided exam questions have been covered, including atomic structure, bonding, molecular geometry, and basic organic chemistry.