Textbook Notes (369,072)
Chemistry (43)
CHEM 112 (33)
Chapter Final

# Chapter 8, 9, 11, 12, 26 and 27 December Exam Review

8 Pages
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Department
Chemistry
Course Code
CHEM 112
Professor
John Carran

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Chapter 8 – Quantum mechanics Electromagnetic radiation – form of energy transmission in which electric and magnetic fields as are propagated as waves through empty space, or through a medium Wave – disturbance that transmits energy through space or a material medium • Energy is discontinuous (energy changes by a jump, or quantum) • Small particles of matter may at times display wave-like properties Photoelectric effect – phenomenon explaining that when light strikes the surface of metals, electrons are ejected Heisenberg uncertainty principle – cannot measure position and momentum accurately , simultaneously Classical mechanics • Objects = particles • Know exact positions and velocities • Energy is continuous Quantum mechanics – accounts for experiments related to atoms, electrons and light that cannot be explained by classical mechanics • Objects = particles or waves • Cannot know exact positions and velocities • Energy is transferred in discrete units  Lyman series, n=1  Balmer series, n=2  Pachen series, n=3 Standing wave – wave that remains in a constant position Assigning Quantum Numbers 1. Principle quantum number, n. Only positive, non-zero number (n=1, 2, 3) - ENERGY LEVEL/MOST PROBABLE DISTANCE OF FINDING e FROM NUCLEUS 2. Orbital angular momentum quantum number, l. Not larger than n-1 (l=0; s, 1; p, 2;d, 3; f) SHAPE OF ORBITAL 3. Magnetic quantum number, m. Negatlve or positive integer, incl 0 and ranging from –l to l (where l is the orbital angular momentum quantum number (m = -l, (-l l+1)….-1, 0, 1, (l-1), +l)) For example, if l = 2, then m can be +/- : -2, -1, 1, 2 and 0. ORIENTATION • ***NOT involved in orbital designation*** • S  spherical, pdumb-bell, d  clover leaf 4. Electron spin, M (+s- 1/2) • All orbitals with same n and l value are in the same sub-shell • Cannot have two electrons in an atom with all 4 quantum numbers the same (Pauli Excl. Princ.) • Only two electrons may occupy the same orbital, and these must have opposing spins • For degenerate orbitals, electrons initially occupy them singly (atom tends to have as many unpaired electrons as possible) Hund’s rule • count the superscripts of the electron configuration of a species to find the atomic number Degenerate – orbitals at the same energy level • For electron configuration, only Cr and Cu are exceptions to the usual filling rule Reference http://answers.yahoo.com/question/index?qid=20090804201613AAGReAq Chapter 9 – Periodic table properties Effective nuclear charge – true nuclear charge minus charge screened out by electrons: Z = Zeff S • Increases across the row, and is constant in each family Diamagnetic atom/ion – all electrons are paired and the individual magnetic effects cancel out Paramagnetic atom/ion – unpaired electrons present and individual magnetic effects don’t cancel out Isoelectric – when atoms or ions have the same number of electrons (same position on periodic table) Metals - more associated to small atomic radius, high ionization energy and negative electron affinities Nonmetals -more associated to greater atomic radius and lower ionization energy • If two ions have the same electronic configuration, in general, the ion with the higher nuclear charge is smaller • Elements within the same row are more related than those in the same family • Atomic radii: o Covalent radius o Metallic radius o Ionic radius • Ionization energy- energy change for the removal of an electron o inverse of radii • Electron affinity- energy change for the addition of an electron, or enthalpy change as atom in gas phase gains an electron. Chapter 12 – Chapter 10 - Chemical Bonding I • Double/triple bonds are more likely to happen with S, N, C, O, H • Formal charge = valence-lone-(1/2)bonded • Expanded valence – nonmetals in period 3 or later (central atom expand to 10 or 12 electrons) • Coordinate covalent bond – when both electrons in a covalent bond are provided by one of the atoms Hybridization (count regions of high electron density; could be lone pair, double or triple bond) • 4 regions = sp  s + p + p + p • 3 regions = sp  s + p + p •
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