Lewis Structures and Resonance

Drawing Lewis Structures

Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They are essential for understanding molecular geometry, resonance, and reactivity.

• Octet Rule: Most main-group elements tend to form bonds until they are surrounded by eight valence electrons.

• Resonance Structures: When more than one valid Lewis structure can be drawn for a molecule or ion, these are called resonance structures. The actual structure is a resonance hybrid of all possible forms.

• Example: The carbonate ion (CO32−) has three resonance structures, each with a double bond to a different oxygen atom.

AXE Notation and Molecular Geometry

The AXE method is used to predict molecular shapes:

• A: Central atom

• X: Number of atoms bonded to the central atom

• E: Number of lone pairs on the central atom

Example: SO2 is classified as AX2E1 (two bonded atoms, one lone pair).

Molecular Shape and Polarity

VSEPR Theory and Molecular Geometry

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

• Common Geometries:

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

  • Trigonal planar: 120° bond angle (e.g., CH2O)

  • Tetrahedral: 109.5° bond angle (e.g., CH4)

  • Bent: < 120° or < 109.5° (e.g., H2O, SO2)

• Example: The indicated carbon in CH3CONHCH3 is trigonal planar due to three regions of electron density (AX3E0).

Bond Polarity and Molecular Polarity

Bond polarity arises from differences in electronegativity between bonded atoms. The direction of the dipole moment is from the less electronegative atom to the more electronegative atom.

• Electronegativity: A measure of an atom's ability to attract electrons in a bond.

• Polarity of Bonds: Indicated by an arrow pointing toward the more electronegative atom.

• Example: In Si–N, the arrow points toward N (more electronegative).

Intermolecular Forces

Types of Intermolecular Forces

• London Dispersion Forces: Present in all molecules, especially significant in nonpolar molecules.

• Dipole-Dipole Forces: Occur between polar molecules.

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

• Ionic Bonds: Electrostatic attraction between ions in ionic compounds.

Example Table:

Physical Properties and Trends

Boiling Point Trends

Boiling points are influenced by molecular size, shape, and intermolecular forces. Generally, larger molecules and those with stronger intermolecular forces have higher boiling points.

• Example: Straight-chain alkanes have higher boiling points than branched isomers due to greater surface area for intermolecular interactions.

Chemical Reactions and Stoichiometry

Balancing Chemical Equations

Balanced chemical equations have equal numbers of each type of atom on both sides. Coefficients are used to achieve this balance.

• Example: For the combustion of hydrazine:

Precipitation Reactions

When two aqueous solutions are mixed, an insoluble product (precipitate) may form if the product is not soluble in water.

• Example: Mixing AgNO3 and K2CrO4 forms Ag2CrO4 precipitate.

Redox Reactions

Oxidation-reduction (redox) reactions involve the transfer of electrons. Oxidation is the loss of electrons, and reduction is the gain of electrons.

• Oxidizing Agent: Causes oxidation, is itself reduced.

• Reducing Agent: Causes reduction, is itself oxidized.

Stoichiometry and Limiting Reactants

Limiting Reactant

The limiting reactant is the reactant that is completely consumed first, limiting the amount of product formed.

• Example: In the reaction , if there is less O2 than needed for all the H2, O2 is the limiting reactant.

Calculating Masses and Moles

Stoichiometric calculations relate the quantities of reactants and products using balanced chemical equations.

• Mole Concept: 1 mole = particles.

• Molar Mass: The mass of 1 mole of a substance (g/mol).

• Example: To find moles of Fe2O3 in 50.00 g, use .

Thermochemistry

Exothermic and Endothermic Reactions

Reactions can either release energy (exothermic) or absorb energy (endothermic).

• Exothermic: is negative; energy is released.

• Endothermic: is positive; energy is absorbed.

Activation Energy and Reaction Diagrams

The activation energy () is the minimum energy required for a reaction to proceed. Reaction energy diagrams show the energy changes during a reaction.

• Reactants: Starting materials.

• Products: Substances formed.

• Activation Energy: The energy difference between reactants and the peak of the curve.

Summary Table: Types of Intermolecular Forces

Additional info:

• Some questions reference specific diagrams or answer choices; the above notes generalize the concepts for study purposes.

• For detailed practice, refer to sample problems on Lewis structures, VSEPR, stoichiometry, and thermochemistry.