1. Tentative material to be covered for Exam 2 (Wednesday, October 27) Chapter 17Many-Electron Atoms and Chemical Bonding 17.1Many-Electron Atoms and the Periodic Table 17.2Experimental Measures of Orbital Energies 17.3Sizes of Atoms and Ions 17.4Properties of the Chemical Bond 17.5Ionic and Covalent Bonds 17.6Oxidation States and Chemical Bonding Chapter 18Molecular Orbitals, Spectroscopy, and Chemical Bonding 18.1Diatomic Molecules 18.2Polyatomic Molecules 18.3The Conjugation of Bonds and Resonance Structures 18.4The Interaction of Light with Molecules 18.5Atmospheric Chemistry and Air Pollution

3. Quantum mechanics provides an intellectual structure for describing all of the properties of atoms and molecules. For atoms quantum mechanics the concept of orbitals (wavefunctions) provides a description of the energies, the sizes of atoms and the basis for bonding of atoms and the construction of the periodic table. The orbitals for the H atom, which are known precisely, are used as starting approximation for building up the electron configuration of multielectron atoms. Every electron in an atom is assigned four quantum numbers (n, l, m l and m s ) that uniquely define its spatial distrbution and spin state. Thus, we can envision every electrons in terms of a characteristic energy, size, shape, orientation and spin.

4. Properties of electrons in atoms Quantum numbers of electrons Electron configurations Core electrons Valence electrons Energy required to remove an electron Energy required to add an electron Size of atoms

5. Building up the ground state configuration of atoms Every atom possesses the SAME set of available orbitals Every electron of an atom MUST be in one of these orbitals: 1s, 2s, 2p, 3s, 3p, 3d, etc. The energy and size of these orbitals depend on the atom (Z), but the shape and orientation is space of any orbitals of the same l are the same for all atoms. The energy ranking of the orbitals for the representative elements is generally: 1s < 2s < 2p < 3s < 3p. From this point on the next lowest energy orbital may be 4s or 3d, depending on the number of electrons in the neutral atom.

6. The properties of the atoms of the elements vary periodically with the atomic weights of the elements. All chemical and physical properties of the elements depend on their atomic weights and therefore vary periodically with atomic weight. The ground state electron configuration of the atoms of elements vary periodically with the atomic number Z. All chemical and physical properties of the elements that depend on electron configurations vary periodically with atomic number.

7. Ground state electron configuration: Z electrons (Z = atomic number of the atom) are placed seriatim into the orbitals according to the following guidelines. Aufbau principle: electrons go into lowest energy orbitals first. Pauli principle: No more than two electrons in any one orbital. Filled orbitals have spins paired. Hund’s rule: When there are orbitals of equal energy in a subshell to fill, the electrons first go into different orbitals with parallel spins one at a time.

8. Valence electrons, Lewis structures and electronic configurations The valence electrons are electrons in the s and p orbitals: valence electrons = s n p m AtomConfigurationComment 3 Li[He]2sParamagnetic 4 Be[He] 2s 2 Closed shell (diamagnetic) 5 B[He] 2s 2 2p 1 Paramagnetic 6 C[He] 2s 2 2p 2 Paramagnetic 7 N[He] 2s 2 2p 3 Paramagnetic 8 O[He] 2s 2 2p 4 Paramagnetic 9 F[He] 2s 2 2p 5 Paramagnetic 10 Ne [He] 2s 2 2p 6 Closed shell (diamagnetic)

9. Correlation of valence electron and Lewis structures

10. Building up the third row of the periodic table: From Na to Ar AtomConfigurationComment 11 Na[Ne]2sParamagnetic 12 Mg[Ne] 2s 2 Closed shell (diamagnetic) 13 Al[Ne] 2s 2 2p 1 Paramagnetic 14 Si[Ne] 2s 2 2p 2 Paramagnetic 15 P[Ne] 2s 2 2p 3 Paramagnetic 16 S[Ne] 2s 2 2p 4 Paramagnetic 17 Cl[Ne] 2s 2 2p 5 Paramagnetic 18 Ar [Ne] 2s 2 2p 6 Closed shell (diamagnetic)

11. d orbitals From photoelectron spectroscopy, the 3d subshell for elements 21 through 29 (Sc through Cu) lies well above the 3d subshell. However, the energy of the 3d subshell is very close in energy to the 4s subshell: 3p << 3d ~ 4s 1s << 2s < 2p << 3s < 3p < 4s ~ 3d Thus is some cases the specifics of orbital configurations place 3d below 4s and in other cases th 4s is below the 3d.

12. The fourth row of the periodic table AtomConfiguration 19 K 18 [Ar]2s 20 Ca 18 [Ar]2s 2 _________________________________ d orbitals fill up 31 Ga 18 [Ar] 2s 2 2p 1 32 Ge 18 [Ar] 2s 2 2p 2 33 As 18 [Ar] 2s 2 2p 3 34 Se 18 [Ar] 2s 2 2p 4 35 Cl 18 [Ar] 2s 2 2p 5 36 Kr 18 [Ar] 2s 2 2p 6 What about 21 M through 30 M?

13. The electron configurations of the transition elements 21 Sc 18 [Ar]4s 2 3d 22 Ti 18 [Ar]4s 2 3d 2 23 V 18 [Ar]4s 2 3d 3 24 Cr 18 [Ar]4s 2 3d 4 instead 18 [Ar]4s 1 3d 5 25 Mn 18 [Ar]4s 2 3d 5 26 Fe 18 [Ar]4s 2 3d 6 27 Co 18 [Ar]4s 2 3d 7 28 Ni 18 [Ar]4s 2 3d 8 29 Cu 18 [Ar]4s 2 3d 9 instead 18 [Ar]4s 1 3d 10 30 Zn 18 [Ar]4s 2 3d 10 The “surprises” for 24 Cr and 29 Cu are due to ignored electron-electron repulsions. For 24 Cr the stability of half shells trumps one filled subshell and a partially filled subshell. For 29 Cu the stability of a full d shell and half filled 4s subshell trumps a partially filled 3d subshell.

14. Example: IE drops dramatically from He to Li. Why? He = 1s 2 versus Li 1s 2 2s. 2s on average further away from nucleus for same average charge (after screening by 1s 2 ).

15. Compare Be = [He]2s 2 versus B = [He]2s 2 2p Compare N = [He]2s 2 2p x 2p y 2p z versus O = [He]2s 2 2p x 2 2p y 2p z

16. Bond length: the distance between the centers (nucleus) of bonded atoms.

17. Atomic radius: the atomic radius of a neutral atom generally decreases from left to right across a period (larger Z) and increases down a group (increase in n).

18. The electron affinity (EA) of an atom is the energy change which occurs when an atom gains an electron. X(g) + e - Xe - (g) Electron affinities of the representative elements: What are the correlations across and down?

19. Electronegativity: a measure of the power of an atom to attract electrons to itself in a bond. Most electronegative atoms: F > O > Cl >N ~ Br > I