EXAM 2

CHEM 1130-4 Fall 2024 Exam 2 Review Notes

Ch 4. Chemical Bonding and Molecular Geometry

4.1 Ionic Bonding

  • Formation of Ions:

    • Cations: Positively charged ions formed by the loss of electrons.

    • Anions: Negatively charged ions formed by the gain of electrons.

  • Ionic Compounds: Formed by the electrostatic attraction between cations and anions.

  • Charge Prediction:

    • Common metallic elements typically form cations, often with charges equal to their group number (for groups 1, 2, and 13).

    • Common nonmetallic elements usually form anions with charges that can be determined by subtracting eight from their group number (for groups 15 through 17).

  • Electron Configurations: Write configurations using standard notation.

4.2 Covalent Bonding

  • Formation of Covalent Bonds: Covalent bonds form when two atoms share one or more pairs of electrons.

  • Electronegativity Definition: Electronegativity is the ability of an atom to attract shared electrons in a bond. Electronegativity increases across a period and decreases down a group.

  • Polarity of Covalent Bonds:

    • Electrons are shifted towards the more electronegative atom in the bond, resulting in partial charges (dipole moments).

4.3 Chemical Nomenclature

  • Periodic Table for Exam: Will not include element names.

  • Memorization Required: Names and symbols for elements 1-36, and additional elements including:

    • Strontium (Sr)

    • Barium (Ba)

    • Palladium (Pd)

    • Platinum (Pt)

    • Silver (Ag)

    • Gold (Au)

    • Mercury (Hg)

    • Tin (Sn)

    • Lead (Pb)

    • Tellurium (Te)

    • Iodine (I)

  • Naming Conventions:

    • Monoatomic cations and anions—transition metals require oxidation state identifiers.

    • Polyatomic cations: Example NH$_{4}^{+}$, Polyatomic anions: Common anions to know.

    • Ionic Compounds: Use cation names followed by anion names.

    • Binary Molecular Compounds: Know the prefixes for 1-10 for naming.

    • Hydrates: Naming hydrates involves identifying water of crystallization.

4.4 Lewis Symbols and Structures

  • Lewis Symbols: Represent valence electrons for neutral atoms and ions using dots.

  • Lewis Structures: Draw dot structures for simple molecules adhering to bonding rules.

  • Octet Rule Definition: Atoms tend to gain, lose, or share electrons to achieve a full outer electron shell of eight electrons.

4.5 Formal Charges and Resonance

  • Formal Charge Calculation:

    • Formula:
      FC = ext{# of valence } e^- - ( ext{# of bonds}) - ( ext{# of lone pair } e^-)

  • Resonance Concept: Molecules can have multiple valid Lewis structures that differ in electron placements but not in the arrangement of atoms.

    • Formal charge analysis can help predict the most favored structure.

    • Factors for stability:

    • Formal Charges: Prefer structures with better charge distribution.

    • Electronegativity: Prefer structures with negative charges on more electronegative atoms.

    • Octet Rule Availability: Prefer structures satisfying octet for all atoms involved.

4.6 Molecular Structure and Polarity

  • Valence Shell Electron Pair Repulsion (VSEPR) Theory: Predicts molecular geometries based on electron pair repulsions.

  • Concepts of Polarity:

    • Polar covalent bonds arise when there is a significant difference in electronegativity.

    • Molecular polarity is determined by the geometry and bond polarities.

  • Lewis Structure Drawing:

    • Use to analyze electron density regions to determine electronic geometry:

    • 2 regions = linear

    • 3 regions = trigonal planar

    • 4 regions = tetrahedral

    • 5 regions = trigonal bipyramidal

    • 6 regions = octahedral

  • Molecular Geometry Determination Rules:

    • If there are NO lone pairs, geometry matches electronic geometry.

    • If there are lone pairs, adjust the molecular geometry accordingly.

  • Bond Dipole Analysis:

    • If bond dipoles cancel, the molecule is nonpolar.

    • If bond dipoles do NOT cancel, the molecule is polar.

Ch 5. Advanced Theories of Bonding

5.1 Valence Bond Theory

  • Covalent Bond Formation: Bonds are formed through the overlap of atomic orbitals.

  • Bond Types:

    • σ-bond Definition: A sigma bond occurs when the electron density is concentrated along the bond axis.

    • π-bond Definition: A pi bond occurs when electron density is above and below the bond axis.

5.2 Hybrid Atomic Orbitals

  • Hybridization Concept: Atomic orbitals can mix to form new, hybridized orbitals suited for bonding.

  • Associated Hybrid Orbitals Based on Geometry:

    • Linear: 2 hybrid orbitals

    • Trigonal planar: 3 hybrid orbitals

    • Tetrahedral: 4 hybrid orbitals

    • Trigonal bipyramidal: 5 hybrid orbitals

    • Octahedral: 6 hybrid orbitals

  • Hybrid System Explanation: Identify mixed atomic orbitals and the unhybridized orbitals remaining.

5.3 Multiple Bonds

  • Multiple Covalent Bonding: Involves the overlap of atomic orbitals to form bonding interactions beyond single bonds.

  • Relation to Resonance and π-bonding: Resonance structures can exhibit different π-bond distributions.

5.4 Molecular Orbital Theory

  • Molecular Orbitals Derivation: Molecular orbitals (MOs) formed from atomic orbitals can accommodate electrons.

  • Molecular Orbital Formation Quantities:

    • Total MOs form equals the summation of atomic orbitals combined.

    • Electrons are placed in MOs according to increasing energy levels.

  • Bonding vs. Antibonding Traits:

    • Bonding orbitals stabilize a molecule, while antibonding orbitals destabilize.

  • Bond Order Calculation:
    ext{Bond order} = \frac{\text{# of e in bonding MO} - \text{# of e in antibonding MO}}{2}

  • Diatomic Molecule Configurations Relation: Electrons contribute to molecular stability and magnetic properties based on unpaired electrons.

5.5 Stability and Magnetism Relation:**

  • If bond order > 0, the molecule is more stable than its separate atoms.

  • If bond order = 0, the molecule is not stable.

  • Unpaired electrons indicate a paramagnetic nature; all paired electrons indicate diamagnetism.

Ch 6. Composition of Substances and Solutions

6.1 Formula Mass

  • Calculation of Formula Masses and Molar Masses:

    • Formula mass is the sum of atomic weights (in amu) of all atoms in a compound.

6.2 Determining Empirical and Molecular Formulas

  • Percentage Composition Calculation: From a given formula to determine the % of each element in that compound.

  • Empirical Formula Calculation:

    • Use a 100 g sample, convert % to grams, and then grams to moles.

    • Assemble a pseudo-formula from moles of each element, converting to the simplest whole number ratio.

  • Empirical Formulas from Combustion Data:

    • Convert mass of combustion products to moles, trace back elemental moles for empirical derivation.

6.3 Molarity

  • Solution Properties:

    • Solute: the substance dissolved.

    • Solvent: the medium that dissolves the solute.

    • Concentration is defined as the amount of solute present in a given volume of solution.

  • Calculating Molarity:

    • Units are moles per liter (M), requires unit conversions.

  • Dilution Calculations:

    • Described by dilution equation:
      C<em>1V</em>1=C<em>2V</em>2C<em>1V</em>1 = C<em>2V</em>2

6.4 Other Units for Solution Concentrations

  • Concentration Units:

    • Mass percentage: mass of solute per 100g of solution

    • Volume percentage: volume of solute per 100mL of solution

    • Mass-volume percentage: mass of solute per volume of solution (e.g., g/mL)

    • Parts-per-million (ppm): mass of solute relative to total mass (1ppm = 1 mg/L)

    • Parts-per-billion (ppb): similar to ppm, for billionth ratios.

  • Perform Computations: Relate volume/mass measures of solutions to concentrations and conversions across units.

Ch 7. Stoichiometry of Chemical Reactions

7.1 Writing and Balancing Chemical Equations

  • Chemical Equations Derivation: Translate narratives into chemical equations, keeping nomenclature rules in mind.

  • Equation Formats:

    • Molecular vs. total ionic formats are different in that molecular represents whole molecules while total ionic shows all ions present.

    • Net ionic format excludes spectator ions, which are ions that do not participate in the reaction.

7.2 Classifying Chemical Reactions

  • Common Reaction Types:

    • Precipitation Reaction: When two solutions combine to form an insoluble solid.

    • Acid-Base Reaction: A reaction where an acid reacts with a base to form water and a salt.

    • Redox Reaction: Involves the transfer of electrons between substances.

  • Ionic Compounds Solubility Trends:

    • Soluble: Compounds with group 1 cations, NH4$^+$, NO3$^-$, etc.

    • Insoluble: OH$^-$, CO3$^{2-}$, etc., unless attached to certain cations.

  • Oxidation States in Redox:

    • Oxidation: Loss of electrons.

    • Reduction: Gain of electrons.

7.3 Reaction Stoichiometry

  • Stoichiometry Concept: Derived from balanced equations providing quantitative relationships among reactants and products.

  • Calculating Stoichiometry: Relate amounts of reactants/products using molar conversions.

7.4 Reaction Yields

  • Theoretical Yield: Maximum amount of product formed from a reaction under ideal conditions.

  • Limiting Reactants: The reactant that is entirely consumed first, limiting product formation.

  • Percent Yield Calculation: It is determined using the formula:
    extPercentYield=Actual YieldTheoretical Yield×100ext{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100

7.5 Quantitative Chemical Analysis

  • Titrations and Gravimetric Analysis:

    • Types of reactions for titrations typically fall under neutralization reactions.

    • Gravimetric analysis is based on precipitate formation and is used to determine mass.

  • Stoichiometric Calculations: Use data from titrations and gravimetric analysis to calculate amounts of reactants and products involved.

  • Combustion Analysis Data for Empirical Formula: Utilize combustion products to backtrack to find compounds’ empirical formulas, also applicable in section 6.2.