CHEM 127 Final Study Guide PDF

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CHEM 127 final review ...

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CHEM 127 Final Study Guide Module 1 Mathematical Toolbox (Chapter 1) • Dimensional analysis (1.4) and important units (density, molarity, molar mass, Joules, Pressures: atm, torr, mmHg, Pa) • SI prefixes (1.4) • Significant figures and uncertainty in measurements (1.5) Atomic Structure (Chapters 2 & 3) • Subatomic particles (2.3) • Atoms/elements (2.5, 2.6) • The mole and molar mass (3.1) • Isotopes and relative abundance (2.5) Interactions with Light and the Bohr Model (Chapter 6) • Wavelike nature of light (6.1) • Particulate nature of light; photoelectric effect (6.1) • Line Spectra (6.1) • Atomic Emission/Absorption (6.2 and tools of the laboratory) Quantum Mechanical Model of the Atom (Chapter 6) • Wave-particle duality, uncertainty principle (6.3) • Schrödinger equation: wavefunctions (Ψ), probability distributions (Ψ2) (6.3) • Orbitals and quantum numbers (6.3) Electronic Structure à Properties (Chapter 8) • Electron Configurations (6.4) • Trends: Zeff, atomic/ionic size, ionization energy, electron affinity (6.5) In order to do well on the final exam, you should master the following learning objectives: • Correctly predict the relative ordering for ionization energy, electron affinity, atomic radius, and ionic radius of a series of elements using effective

nuclear charge and atomic shell arguments • Correctly determine ground state electron configurations for elements • Identify different isotopes and correctly determine the number of protons, neutrons, and electrons • Determine the average molar mass of an element from isotopic abundances • Correctly identify quantum numbers associated with an electron in a particular atomic orbital • Explain why atoms emit light at only specific wavelengths in terms of quantized electron energies • Use the Rydberg equation to determine the energy associated with a particular electron transition in a hydrogen atom • Use the relationships between the wavelength, frequency, and energy of a photon to determine the energy, wavelength, or frequency of a photon emitted or absorbed by a hydrogen atom • Associate increases and decreases of electron energy with absorption and emission processes • Use their understanding of the wave nature of light to describe electromagnetic radiation in terms of its wavelength, frequency, and color (if visible). • Use their understanding of the particle nature of light and light quanta or photons to relate the energy of a photon of light to its wavelength or frequency • Use an absorption spectrum to determine the wavelength of maximum absorbance and predict the color of the material absorbing light • Explain how absorbance of light, concentration, and molar absorptivity are related

Module 2 Ionic Bonding (Chapters 2, 3, 7, and 10) • Nomenclature (2.6-2.7) • Ionic bonding model – Lattice energy (trends based on charge density; no Born-Haber cycle) (7.1, 7.5) • Empirical formulas (2.4, 3.2) • Basic crystal structure (unit cell stoichiometry) (10.6) Covalent Bonding and Molecular Structure (Chapters 2, 3, 7, 8, 20) • Nomenclature (2.6-2.7) • Covalent bonding model (7.2) • Empirical and Molecular Formulas (2.4, 3.2)

• Lewis diagrams (7.3) • Line-angle drawings (20.1) • Isomers (2.4 plus additional material from activity) • Resonance: localization vs delocalization (7.4) • Functional groups: alakane, akene, alkyne, alcohol, ether, aldehyde, ketone, carboxylic acid, ester • VSEPR and molecular polarity (7.6) • Hybrid orbitals (up to sp3), σ and π bonding (8.1-8.3) • Bond length, order and strength (7.5) Intermolecular Forces (Chapter 10) • Types of intermolecular forces (10.1) • Energy associated with IMFs (10.1) In order to excel on the final exam, you should master the following learning objectives and understand how these concepts connect with each other: • Understand what a mole quantity is, and how to use it -- grams molesatoms/molecules • Predict which elements will form cations and anions, their charge, and electron configurations, and how this relates to ionic bonding • Predict bonding behaviors of Hydrogen, Carbon, Nitrogen and Oxygen using valence electrons • Recognize and explain bonding patterns for simple molecules • Describe the Octet Rule • Sketch good Lewis Structures from chemical formulae • Create multiple good Lewis Structures of different isomers from chemical formulae • Identify functional groups in a Lewis Structure • Create good Lewis Structures with specific functional groups from chemical formulae • Predict Electronic and Molecular Geometry by identifying the number and type of electron groups around a central atom • Translate chemical formulae into three-dimensional molecular representations using Lewis Structures and VSEPR • Explain relative strengths of intermolecular forces in molecules based on Lewis structures and 3D models • Describe the relationships between elution orders (in GC), boiling points, and intermolecular forces

• Explain relative boiling points using intermolecular forces

Module 3 Types of Reactions and Reaction Stoichiometry (Chapter 4) • Types of reactions: precipitation, redox (no half-reaction balancing), acid-base (4.1 – 4.5) • Writing equations: molecular and ionic (4.1) • Stoichiometry (4.3) • Limiting Reactant, percent yield (4.4) Solution Stoichiometry (Chapters 3 and 4) • Molarity (3.3) • Solution reaction stoichiometry (4.5) Thermochemical Stoichiometry (Chapters 5 and 7, 8) • ΔHrxn stoichiometry (skip calorimetry) (5.3) • Bond enathlpies in reactions (bonds broken – bonds formed) (7.4, 8.1) Gas Phase Stoichiometry (Chapter 9) • Gas Phase Stoichiometry (9.3) • Ideal Gas Law in stoichiometry (9.3) In order excel on the final exam, you should master the following learning objectives and understand how these concepts connect with each other: • Identify common macroscopic signs that a chemical reaction has occurred • Use patterns in chemical reactivity to predict the products and write balanced chemical equations • Use stoichiometry to determine the concentration of an unknown solute in an aqueous solution via titration • Use stoichiometry to determine the empirical and molecular formulae of an unknown • Perform acid-base titrations • Use reaction stoichiometry and quantitative measurements to determine the composition of a metal carbonate/metal bicarbonate mixture • Apply algebraic concepts to reaction stoichiometry in cases where there are multiple unknown quantities • Use the ideal gas law to quantitatively track reactions that involve gases...