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Review of Atomic Orbitals 1923 – Electron has the properties of a particle and a w a v e (constructive/destructive interference) 1926 – Quantum Mechanics • Schrödinger solves a “wave equation” for the total energy of 1 electron and 1 proton. • A series of “wave functions” results. These correspond to the various states of an electron. From wave functions (ψ), we get 1) The energy associated with an electronic state 2) The relative probability of finding an electron in a location in space (ψ 2) E 3p (3) 3s 2p (3) 2s 1s f and d levels are not very important in organic chemistry Shapes of Atomic Orbitals a “wave” + + + – – Signs are not related to electric charge or probability –/–, +/+ interfere constructively, +/– interfere destructively 1s-orbital 2s-orbital nucleus node (–) (+) • (+) • y (+) – – + z (–) 2py-orbital x 2px-orbital + 2pz-orbital 1s, no nodes; 2s & p, 1 node; 3 s & p, 2 nodes, etc.… “Orbitals” are regions in space where the probability of finding an electron is high (ψ 2 = 95%). Orbitals are not “real”! Filling Atomic Orbitals 3p 3s 2p degenerate orbitals equal in energy 2s 1s Orbitals are filled in this order with the following rules. 1) Aufbau principle place electrons in lowest level, work up max 2/per orbital 2) 3) Pauli exclusion principle. Electrons have two possible spin states. To be in an orbital together, they must be paired (opposite) or or paired unpaired Hundʼs Rule Add one electron to each orbital with spins unpaired, when each has one, double up and pair. Why? Electrons repel each other! C N O F 4 3 2 1 Ne full Shapes of Molecules by VSEPR Methane Ammonia Boron Trifluoride Theory