Chemistry Chapters 3-5 Exam Review
Chapter 3 Notes: Atomic Structure and Quantum Chemistry
Learning Objectives
Electromagnetic Radiation:
Convert between energy (E), wavelength (\lambda), and frequency (\nu) of electromagnetic radiation. The relationship is given by E = h \nu = \frac{hc}{\lambda}, where h is Planck's constant (6.626 \times 10^{-34} \text{ J} \cdot \text{s}) and c is the speed of light (2.998 \times 10^8 \text{ m/s}).
Link these values to their appropriate regions within the electromagnetic spectrum (e.g., radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, gamma rays).
Quantum Numbers and Orbitals:
Assign quantum numbers (n, l, ml, ms) to atomic orbitals.
Use these quantum numbers to describe the size (n), energy (n for hydrogenic atoms, n and l for multi-electron atoms), and orientation (m_l) of orbitals.
Principal Quantum Number (n): Describes the energy level and size of the orbital. Can be any positive integer (1, 2, 3, …).
Azimuthal (Angular Momentum) Quantum Number (l): Describes the shape of the orbital. Can range from 0 to n-1. l=0 corresponds to an s orbital (spherical), l=1 to a p orbital (dumbbell shape), l=2 to a d orbital (more complex shapes), and so on.
Magnetic Quantum Number (ml): Describes the orientation of the orbital in space. Can range from -l to +l in integer steps. For l=1 (p orbitals), ml can be -1, 0, +1, corresponding to px, py, p_z orientations.
Spin Quantum Number (m_s): Describes the intrinsic angular momentum (spin) of an electron. Can only be +\frac{1}{2} or -\frac{1}{2}.
Electronic Configurations and Orbital Diagrams:
Apply the Aufbau principle (electrons fill lowest energy orbitals first).
Apply Hund's Rule (when degenerate orbitals are available, electrons fill them singly with parallel spins before pairing up).
Write electronic configurations (e.g., 1s^2 2s^2 2p^6) and orbital diagrams (showing individual spin states in boxes/lines) for atoms and monoatomic ions.
Periodic Trends:
Relate Ionization Energies (energy required to remove an electron from a gaseous atom or ion) to the element's position on the periodic table. Generally increases across a period and decreases down a group.
Relate Electron Affinities (energy change when an electron is added to a gaseous atom) to the element's position on the periodic table. Generally becomes more negative (more exothermic) across a period and less negative down a group (with exceptions).
Suggested Practice & Review Materials
Animations: Review the worked problems within the animation links found on the Chapter 3 Preview page, especially those assigned for Preclass Quiz.
Textbook Sample Exercises: Rework examples from paragraphs 3.1, 3.2, 3.8, 3.9, 3.10, 3.11, 3.12, 3.13.
End-of-Chapter Problems (Textbook): For additional practice: 3.10, 3.75, 3.79, 3.89, 3.91, 3.93, 3.94, 3.99, 3.107, 3.109, 3.115.
Class Slide Shells: Re-solve and review all examples within the Chapter 3 slide shells (available in both blank and annotated/answered versions).
Recitation Handout for Chapter 3:
Contains an outline of all topics and examples.
Answer Key (AK) for these examples is provided in the Exam Module.
It is crucial to review the content and attempt to solve the handout examples independently first, focusing on setting up the problems.
Use the handout to effectively map concepts, definitions, and equations to problem-solving strategies.
Old Exam Sample: Available in the Exam Review Module in Canvas.
Chapter 4 Notes: Bonding Connectivities (Lewis Structures) and Types of Bonding
Core Concepts (Continuing from Ramp-up Modules 6 & 7)
Lewis Structures: Enhanced building techniques for correct structures.
Resonance: Drawing multiple valid Lewis structures when a single structure cannot adequately represent bonding.
Formal Charge: A tool to evaluate the most plausible Lewis structure; FC = (\text{valence electrons}) - (\text{non-bonding electrons}) - \frac{1}{2}(\text{bonding electrons}).
Extended Octet: When central atoms in period 3 or below can accommodate more than 8 valence electrons.
Bond Order: The number of chemical bonds between a pair of atoms. A bond order of 1 is a single bond, 2 is a double bond, and 3 is a triple bond. Affects bond energy and length.
Odd Electron Species (Radicals): Molecules or ions with an odd number of valence electrons, resulting in one or more unpaired electrons.
Ionic Bonds: Refinement of understanding.
Ionic Lattice Energy: The energy required to separate one mole of an ionic solid into its gaseous constituent ions. Comparisons of lattice energy are crucial for understanding the strength of ionic bonds and melting points.
Bond Polarity and Electronegativity:
Electronegativity (X): A measure of the tendency of an atom to attract a bonding pair of electrons.
Electronegativity Difference (\Delta X): The difference in electronegativity between two bonded atoms. Used to predict the polarity of covalent bonds:
\Delta X \approx 0: Purely covalent (nonpolar)
0 < \Delta X \le 0.4: Nonpolar covalent
0.4 < \Delta X \le 1.7: Polar covalent
\Delta X > 1.7: Ionic
Molecular Polarity: Build foundations for understanding molecular polarity (Chapter 5), including its role in the Greenhouse Effect.
Learning Outcomes
LO1 - Bonding Types: Describe the similarities and differences between covalent, ionic, and metallic bonds.
LO2 - Ionic Compound Strengths: Calculate and compare the relative strengths of ion-ion attractions in ionic compounds. (See Sample 4.1 in textbook).
LO3 - Covalent Bond Polarity: Predict the polarity of covalent bonds based on the electronegativity difference (\Delta X) between bonded atoms. (See Sample 4.2 in textbook).
LO4 - Naming Compounds & Formulas: Expand on Chapter 2 concepts to name molecular and ionic compounds and write their correct formulas. (See samples in textbook paragraphs 4.3 - 4.8).
LO5 - Lewis Structures: Draw Lewis structures for molecular compounds and polyatomic ions. (See samples in textbook paragraphs 4.9 - 4.13).
LO6 - Resonance & Formal Charge: Draw resonance structures and use formal charges to evaluate and finalize the best possible structures. (See samples in textbook paragraphs 4.14 and 4.16 - 4.18).
LO7 - Bond Properties: Describe the relationship between bond order, bond energy, and bond length. Higher bond order leads to higher bond energy and shorter bond length.
LO8 - Greenhouse Effect: Explain how molecules absorb infrared (IR) radiation and contribute to the Greenhouse Effect.
Suggested Practice & Review Materials
Animations: Review worked problems from the Chapter 4 Preview page animation links, especially those assigned for the Preclass Quiz.
Textbook Worked Examples: Rework examples within textbook paragraphs 4.1 - 4.9.
Class Slide Problems: Review and re-solve all problems from the Chapter 4 slide sets (available as blank files and with lecture annotations).
Recitation Handout for Chapter 4:
Contains summaries of concepts and equations from Chapter 4 tied to specific examples.
Answer Key (AK) for these examples is provided in the Exam Module.
Solve these problems to practice:
Setting up and refining Lewis structures via a six-step process.
Using formal charge to evaluate structures.
Drawing resonance structures.
Applying the extended octet rule or addressing exceptions (e.g., cases where octet is not achieved, including odd electron/radical species).
Ranking melting points of ionic solids based on ionic bond strengths.
Naming ionic and covalent compounds.
Determining bond polarity.
Old Exam: Available in the Exam Module.
End-of-Chapter Problems (Textbook): For additional practice on building structures: 4.5, 4.9, 4.33, 4.35, 4.37, 4.77, 4.89, 4.99, 4.101 - 4.105, 4.116, 4.139, 4.145.
Chapter 5 Study Guide: 3-D Molecular Structures, Molecular Polarity, Hybridization, Chirality
Focus for Exam #1
This module explores the three-dimensional (3-D) shapes of molecules using optimized Lewis structures. For Exam #1, the focus will be on the Valence Shell Electron Pair Repulsion (VSEPR) Theory and Molecular Polarity.
Sections Covered for Exam #1:
5.2 Valence Shell Electron Pair Repulsion Theory (VSEPR)
5.3 Polar Bonds and Molecules
Topics NOT on Exam #1 (but will be covered): Hybridization and Chirality.
Learning Objectives
VSEPR Theory:
Assign steric numbers (number of electron domains around the central atom, including bonding pairs and lone pairs) and electron-pair geometry to Lewis structures.
Use VSEPR theory to determine the molecular geometry (shape of the molecule considering only atom positions) of a molecule based on its electron-pair geometry and the presence of lone pairs.
Predict bond angles in molecules with varying molecular shapes (e.g., tetrahedral, trigonal planar, linear, bent, trigonal pyramidal).
Molecular Polarity:
Identify bond dipoles (vectors pointing from less electronegative to more electronegative atom) within a molecule.
Estimate the relative polarity of compounds by considering the magnitude and orientation of individual bond dipoles and determining if they cancel out or result in a net dipole moment.
Suggested Practice & Review Materials
Animations (Chapter 5 Module):
Molecular Shape and Polarity: Review these links thoroughly.
VSEPR Theory and Basic Shapes: Focus specifically on knowing steric numbers 2-4 (linear, trigonal planar, tetrahedral) without any external support. For steric numbers 5 (trigonal bipyramidal) and 6 (octahedral), focus only on structures where all positions are occupied by bonding atoms (no lone pairs).
VSEPR Theory and the Effect of Lone Pairs: Understand how lone pairs influence molecular geometry and bond angles.
Textbook (Additional Readings): Review worked examples within textbook paragraphs 5.1 - 5.6.
Recitation Handout for Chapters 3-5:
Answer Key (AK) is provided in the Exam Module.
It's important to review content and attempt examples independently first, focusing on problem setup.
Use the booklet to map concepts, definitions, and equations to problems.
Class Slide Problems: Review and re-solve all problems from the Chapter 5 slide sets (available in both blank and annotated versions).
Recitation Booklet Examples: Use all examples and worked problems to practice applying learning objectives and building a concept map for:
Setting up and refining simple Lewis structures to determine the correct 3-D molecular geometry.
Determining molecular polarity by placing dipole arrows on bonds within a correct 3-D structure, assessing whether they cancel or not, and concluding if the molecule is polar or nonpolar.
Additional Practice: The animation problems, slide problems, and recitation booklet/quiz problems provide a comprehensive set for independent re-solving to ensure a clear understanding of building correct 3-D shapes and determining molecular polarity.
Old Exam Sample: Found in the Exam Review Module.
End-of-Chapter Problems (Textbook): For further practice in building structures: 5.1, 5.3, 5.19, 5.29, 5.43, 5.45, 5.53.