Tentative material to be covered for Exam 2 (Wednesday, October 27)

Chapter 17 Many-Electron Atoms and Chemical Bonding

17.1 Many-Electron Atoms and the Periodic Table17.2 Experimental Measures of Orbital Energies17.3 Sizes of Atoms and Ions17.4 Properties of the Chemical Bond17.5 Ionic and Covalent Bonds17.6 Oxidation States and Chemical Bonding

Chapter 18 Molecular Orbitals, Spectroscopy, and Chemical Bonding

18.1 Diatomic Molecules18.2 Polyatomic Molecules18.3 The Conjugation of Bonds and Resonance Structures18.4 The Interaction of Light with Molecules18.5 Atmospheric Chemistry and Air Pollution

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, ml and ms) 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.

Properties of electrons in atoms

Quantum numbers of electronsElectron configurationsCore electronsValence electronsEnergy required to remove an electronEnergy required to add an electronSize of atoms

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.

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.

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.

Valence electrons, Lewis structures and electronic configurations

The valence electrons are electrons in the s and p orbitals: valence electrons = snpm

Atom Configuration Comment3Li [He]2s Paramagnetic4Be [He] 2s2 Closed shell (diamagnetic)5B [He] 2s22p1 Paramagnetic6C [He] 2s22p2 Paramagnetic7N [He] 2s22p3 Paramagnetic8O [He] 2s22p4 Paramagnetic9F [He] 2s22p5 Paramagnetic10Ne [He] 2s22p6 Closed shell (diamagnetic)

Correlation of valence electron and Lewis structures

N [He]2s22px2py2pz

O [He]2s22px22py2pz

F [He]2s22px22py

22pz

Ne [He]2s22px22py

22pz2

Filled shell

Building up the third row of the periodic table:

From Na to Ar

Atom Configuration Comment

11Na [Ne]2s Paramagnetic12Mg [Ne] 2s2 Closed shell (diamagnetic)13Al [Ne] 2s22p1 Paramagnetic14Si [Ne] 2s22p2 Paramagnetic15P [Ne] 2s22p3 Paramagnetic16S [Ne] 2s22p4 Paramagnetic17Cl [Ne] 2s22p5 Paramagnetic18Ar [Ne] 2s22p6 Closed shell (diamagnetic)

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.

The fourth row of the periodic table

Atom Configuration

19K 18[Ar]2s20Ca 18[Ar]2s2

_________________________________ d orbitals fill up31Ga 18[Ar] 2s22p1

32Ge 18[Ar] 2s22p2

33As 18[Ar] 2s22p3

34Se 18[Ar] 2s22p4

35Cl 18[Ar] 2s22p5 36Kr 18[Ar] 2s22p6

What about 21M through 30M?

The electron configurations of the transition elements

21Sc 18[Ar]4s23d22Ti 18[Ar]4s23d2

23V 18[Ar]4s23d3

24Cr 18[Ar]4s23d4 instead 18[Ar]4s13d5

25Mn 18[Ar]4s23d5

26Fe 18[Ar]4s23d6

27Co 18[Ar]4s23d7

28Ni 18[Ar]4s23d8

29Cu 18[Ar]4s23d9 instead 18[Ar]4s13d10

30Zn 18[Ar]4s23d10

The “surprises” for 24Cr and 29Cu are due to ignored electron-electron repulsions.For 24Cr the stability of half shells trumps one filled subshell and a partially filled subshell.For 29Cu the stability of a full d shell and half filled 4s subshell trumps a partially filled 3d subshell.

Example: IE drops dramatically from He to Li. Why? He = 1s2 versus Li 1s22s. 2s on average further away from nucleus for same average charge (after screening by 1s2).

Compare Be = [He]2s2 versus B = [He]2s22pCompare N = [He]2s22px2py2pz versus O = [He]2s22px

22py2pz

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

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).

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?

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