Inorganic Chemistry Notes Ch2

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Chapter 2: Atomic Structure Quantum Numbers n: principal, energy, size of orbital, positive integers allowed l: angular momentum, orbital shape, integers from 0 to n-1 allowed ml: magnetic moment, orientation in space of orbital, integers from -l to +l allowed ms: spin, shows electrons in orbital, ± ½ allowed Orbitals p orbital l = 1 All have a specific orientation in space (x, y, z) Electron spends equal amount of time in both lobes, nucleus at node Shading shows positive and negative values of wavefunction (assymetrical with respect to inversion d orbital l = 2 d orbitals have 5 orientations in space dyx, dxy, dxz, dz2 Symmetrical with respect to inversion Filling out orbitals Each orbital has a max of 2 electrons Pauli exclusion principle: no 2 electrons in an atom can have the same 4 quantum numbers Up spin filled first Takes less energy Each orbital has its up-spin filled first so as to minimize Coulombic repulsion Hund’s Rules of Maximum of Multiplicity Multiplicity = unpaired electrons + 1 Higher is better Coulombic energy of repulsion (ΠC) Two electrons in the same orbital have higher energy than two electrons in different orbitals Coulombic energy of exchange (Πe) Possible exchanges between two electrons with the same energy and spin (more available exchanges is better) Repulsion energy is destabilizing (ΠC>0), exchange energy is stabilizing (Πe<0) Π = ΠC + Πe Electrons are placed in orbitals to give the lowest total energy to the atom Orbitals filled using Aufbau Chromium, Molybdenum, Copper, Silver, and Gold are all exceptions to the Aufbau principle Cr: [Ar] 4s1 3d5 Many times, half-filled and completely filled orbitals are unusually stable, so sigma electrons can rarely end up promoted to pi electrons Types of electrons Valence: outer electrons responsible for chemistry Transition metals: Any

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