Tentative material to be covered for Exam 2
(Wednesday, October 27)
Chapter 17
Many-Electron Atoms and Chemical Bonding
17.1
17.2
17.3
17.4
17.5
17.6
Many-Electron Atoms and the Periodic Table
Experimental Measures of Orbital Energies
Sizes of Atoms and Ions
Properties of the Chemical Bond
Ionic and Covalent Bonds
Oxidation States and Chemical Bonding
Chapter 18
Molecular Orbitals, Spectroscopy, and Chemical Bonding
18.1
18.2
18.3
18.4
18.5
Diatomic Molecules
Polyatomic Molecules
The Conjugation of Bonds and Resonance Structures
The Interaction of Light with Molecules
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 electrons
Electron configurations
Core electrons
Valence electrons
Energy required to remove an electron
Energy required to add an electron
Size 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
3Li
4Be
5B
6C
7N
8O
9F
10Ne
Configuration
[He]2s
[He] 2s2
[He] 2s22p1
[He] 2s22p2
[He] 2s22p3
[He] 2s22p4
[He] 2s22p5
[He] 2s22p6
Comment
Paramagnetic
Closed shell (diamagnetic)
Paramagnetic
Paramagnetic
Paramagnetic
Paramagnetic
Paramagnetic
Closed shell (diamagnetic)
Correlation of valence electron and Lewis structures
N
[He]2s22px2py2pz
O
[He]2s22px22py2pz
F
[He]2s22px22py22pz
Ne
[He]2s22px22py22pz2
Filled shell
Building up the third row of the periodic table:
From Na to Ar
Atom
Configuration
Comment
11Na
[Ne]2s
[Ne] 2s2
[Ne] 2s22p1
[Ne] 2s22p2
[Ne] 2s22p3
[Ne] 2s22p4
[Ne] 2s22p5
[Ne] 2s22p6
Paramagnetic
Closed shell (diamagnetic)
Paramagnetic
Paramagnetic
Paramagnetic
Paramagnetic
Paramagnetic
Closed shell (diamagnetic)
12Mg
13Al
14Si
15P
16S
17Cl
18Ar
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]2s
20Ca
18[Ar]2s2
_________________________________
31Ga
18[Ar]
33As
18[Ar]
32Ge
34Se
35Cl
36Kr
18[Ar]
18[Ar]
18[Ar]
18[Ar]
2s22p1
2s22p2
2s22p3
2s22p4
2s22p5
2s22p6
What about
21M
d orbitals fill up
through
30M?
The electron configurations of the transition elements
21Sc
22Ti
23V
24Cr
25Mn
26Fe
27Co
28Ni
29Cu
30Zn
18[Ar]4s23d
18[Ar]4s23d2
18[Ar]4s23d3
18[Ar]4s23d4
18[Ar]4s23d5
instead 18[Ar]4s13d5
18[Ar]4s23d6
18[Ar]4s23d7
18[Ar]4s23d8
18[Ar]4s23d9
18[Ar]4s23d10
instead 18[Ar]4s13d10
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]2s22p
Compare N = [He]2s22px2py2pz versus
O = [He]2s22px22py2pz
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) + eXe- (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