Exam 3 Review Guide PDF

Learning Objectives =================== Chapter 6 --------- Explain the basic behavior of waves, including travelling waves and standing waves Describe the wave nature of light Use appropriate equations to calculate related light-wave properties such as frequency, wavelength, and energy...

Learning Objectives =================== Chapter 6 --------- Explain the basic behavior of waves, including travelling waves and standing waves Describe the wave nature of light Use appropriate equations to calculate related light-wave properties such as frequency, wavelength, and energy Distinguish between line and continuous emission spectra Describe the particle nature of light Describe the Bohr model of the hydrogen atom Use the Rydberg equation to calculate energies of light emitted or absorbed by hydrogen atoms Extend the concept of wave--particle duality that was observed in electromagnetic radiation to matter as well Understand the general idea of the quantum mechanical description of electrons in an atom, and that it uses the notion of three-dimensional wave functions, or orbitals, that define the distribution of probability to find an electron in a particular part of space List and describe traits of the four quantum numbers that form the basis for completely specifying the state of an electron in an atom Derive the predicted ground-state electron configurations of atoms Identify and explain exceptions to predicted electron configurations for atoms and ions Relate electron configurations to element classifications in the periodic table Describe and explain the observed trends in atomic size, ionization energy, and electron affinity of the elements Chapter 7 --------- Explain the formation of cations, anions, and ionic compounds Predict the charge of common metallic and nonmetallic elements, and write their electron configurations Describe the formation of covalent bonds Define electronegativity and assess the polarity of covalent bonds Write Lewis symbols for neutral atoms and ions Draw Lewis structures depicting the bonding in simple molecules Compute formal charges for atoms in any Lewis structure Use formal charges to identify the most reasonable Lewis structure for a given molecule Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule Describe the energetics of covalent and ionic bond formation and breakage Use the Born-Haber cycle to compute lattice energies for ionic compounds Use average covalent bond energies to estimate enthalpies of reaction Predict the structures of small molecules using valence shell electron pair repulsion (VSEPR) theory Explain the concepts of polar covalent bonds and molecular polarity Assess the polarity of a molecule based on its bonding and structure Chapter 8 --------- Describe the formation of covalent bonds in terms of atomic orbital overlap Define and give examples of σ and π bonds Explain the concept of atomic orbital hybridization Determine the hybrid orbitals associated with various molecular geometries Describe multiple covalent bonding in terms of atomic orbital overlap Relate the concept of resonance to π-bonding and electron delocalization Outline the basic quantum-mechanical approach to deriving molecular orbitals from atomic orbitals Describe traits of bonding and antibonding molecular orbitals Calculate bond orders based on molecular electron configurations Write molecular electron configurations for first- and second-row diatomic molecules Relate these electron configurations to the molecules' stabilities and magnetic properties Topics Covered ============== Chapter 6 --------- Electromagnetic radiation Waves Wavelength, frequency, amplitude Quantization and nodes Energy, photons, photoelectric effect Wave-particle duality Absorption and emission spectra Bohr model Ground and excited states Energy and wavelength of electron transitions Quantum numbers De Broglie wavelengths and Heisenberg uncertainty principle Wavefunctions and quantum mechanics -- where do we find an electron Quantum numbers -- n, l, ml, and ms Orbitals -- s, p, d, f, etc. Pauli exclusion principle Electron configurations and orbital diagrams Aufbau principle and Hund's rule Core vs valence electrons Covalent and ionic radius Ionization energy and electron affinity Effective nuclear charge (Zeff) Chapter 7 --------- Ionic bonds and formation of ionic compounds Electronic structure of ions Covalent bonds and covalent molecules Bond length Pure vs polar covalent bonds (polarity) Electronegativity and bond polarity Lewis symbols and what they convey Ionic bonding using Lewis structures Lone pairs vs bonding pairs Octet rule and deviations from it Single vs double vs triple bonds Formal charge and resonance structures Formal charge = valence electrons -- (lone pair e- + \# of bonds) Molecular structure prediction using formal charge (finding the "preferred" structure) Covalent bond strength and length Ionic bond strength and lattice energy Born Haber Cycle (calculating lattice energy or associated values) Molecular structure and the polarity of molecules Factors in the electronegativity differences and shape of molecules VSEPR theory Different groups have different size and space requirements around a central molecule Electron-pair vs molecular geometry -- can be the same, but different if there are lone pairs Equatorial vs axial positions Dipole moment -- a measure of bond polarity Chapter 8 --------- Valence bond theory Overlap of half-filled atomic orbitals to give molecular orbitals Sigma vs pi bonds Hybrid atomic orbitals Go through hybridization to form new hybrid orbitals sp, sp2, sp3, sp3d, sp3d2, etc. hybridization Role of hybridization in multiple bond -- unhybridized orbitals can form pi bonds Molecular orbital theory Also explains the magnetic behavior of molecules Atomic orbitals combine to form bonding and antibonding molecular orbitals Molecular orbital diagram Bond order calculations Need to compare the number of electrons in bonding vs antibonding orbitals