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Unformatted text preview: Molecularity Molecularity CH1010
Chapter 8 Part II Electron Configuration and Chemical Periodicity James P. Dittami Orbitals for the Hydrogen Atom Orbitals for the Hydrogen Atom Described by the quantum numbers that arise from the Schrödinger Equation n Principal Quantum no. 1, 2, 3 … l Angular momentum Quantum no. 0 to n1 Energy and Size ml Magnetic Quantum no. –l to l Shape – s, p, d, f Orientation Fourth Quantum No. ms Fourth Quantum No. Electrons have charge and magnetic moment Conclusion electrons have spin ms the fourth quantum no. with value spin = +1/2 or 1/2. Pauli’s Exclusion Principle Pauli’s Exclusion Principle “In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms)” Since electrons in the same orbital have the same values of n, l, ml, this principle says they must have different values of ms Also, since ms = ± ½ (only two values) there can only be 2 electrons/ orbital Figure 7.18 Orbital Energy Figure 7.18 Orbital Energy Levels for the Hydrogen Atom Polyelectronic Atoms Polyelectronic Atoms Revise our picture of orbital energy levels Electron electron repulsion in the atom Inner electrons shield outer electrons from the nucleus Revisit the energy level diagram for Hydrogen to reflect these considerations For a given value of n not all electrons are of same energy The Orders of the Energies of the Orbitals in The Orders of the Energies of the Orbitals in the First Three Levels of Polyelectronic Atoms Electronic Configurations Orbitals will be “H” like (Aubau Principle) Electrons fill from lowest to highest Energy Electrons fill degenerate orbitals to give maximum no. of unpaired electrons (Hund’s Rule) e.g. px , py , pz . Each orbital accomodates Maximum of 2 electrons (Pauli Exclusion Principle) Assign Electronic Configurations Walk through Periodic Table filling each level one by one. Figure 2.10 The modern periodic table. Assign Electronic Configurations H 1s1 He 1s2 Li 1s2 2s1 Be 1s2 2s2 B 1s2 2s2 2p1 C 1s2 2s2 2p2 where 2px1 2py1 and 2pz0 Assign Electronic Configurations C 1s2 2s2 2p2 where 2px1 2py1 and 2pz0 N 1s2 2s2 2p3 where 2px1 , 2py1 and 2pz1 O 1s2 2s2 2p4 where 2px2 , 2py1 and 2pz1 F 1s2 2s2 2p5 where 2px2 , 2py2 and 2pz1 Ne 1s2 2s2 2p6 where 2px2 , 2py2 and 2pz2 Moving to next Row Reference Last Config Na 1s2 2s2 2p6 3s1 or simply [Ne] 3s1 Inner (Core) Electrons Inner Core electrons are those seen in the previous noble gas and any completed transition metal series the lowest E levels of the atom e.g. Na 1s2 2s2 2p6 3s1 or [Ne] 3s1 Core electrons are 1s2 2s2 2p6 Outer Electrons Outer Electrons are those in the highest Energy Level Highest n value e.g. N 1s2 2s2 2p3 Outer electrons are Outermost principle quantum no. n = 2 Outer electrons are 2s2 2p3 for total of 5 Highest energy Least tightly Bound (furthest from nucleus) Valence Electrons Valence electrons those involved in bonding to form compounds For Main Group Elements Valence e are the Outer Electrons e.g. N 1s2 2s2 2p3 For transition metals Count also the (n1)d electrons e.g. Ti [Ar]4s2 3d2 Valence electrons are 2s2 2p3 for total of 5 Valence electrons are 4s2 3d2 for total of 4 Electronic Configurations Beyond Ar Fill K, Ca then transitionmetals BUT: Next fill transition metals with nd electrons Always fill (n+1)s before nd electrons E.g. K 1s2 2s2 2p6 3s2 3p6 4s1 or [Ar] 4s1 E.g. Ca 1s2 2s2 2p6 3s2 3p6 4s2 or [Ar] 4s2 Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 or [Ar]4s2 3d1 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 or [Ar]4s2 3d2 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 or [Ar]4s2 3d3 Figure 8.11 A periodic table of partial ground-state electron configurations Mnemonic for assigning electronic configurations ...
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