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 (will not be covered on exam)
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
New material on course web site:
Practice exams for Chapters 17 and 18 are on line
Answers to Chapter 16 will be distributed by tomorrow
morning
The slide show for Exam 2 has been upgraded. For
those having trouble viewing it, please go to room 211
Havemeyer where the slide show is available from 9-5
A new Powerpoint slide show on material in Chapter 18
has been added
Multielectron atoms:
Every electron in a multielectron atom is assigned
four quantum numbers (n, l, ml and ms) that
uniquely define its chemical and physical properties.
From the use of quantum numbers, we can envision
every electron of a multielectron atom in terms of
a characteristic energy (En), size (r), shape (l),
orientation (ml) and spin (ms).
The highest energy (valence) electrons are of
greatest chemical interest
The electron configurations of multielectron atoms
Row
Configuration
Shorthand
First:
1s2
2[He]
Second:
2s22p6
10[Ne]
Third:
3s23p6
18[Ar]
Fourth:
4s23d104p6
36[Kr]
Fifth:
5s24d105p6
54[Xe]
Perodic Law by mass:
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 vary periodically with
atomic weight.
Periodic Law by electronic structure:
The ground state electron configuration of the atoms of
elements vary periodically with the atomic number Z.
The chemical and physical properties of the elements
depend on the electron configurations of the atoms and
vary periodically with atomic number.
Bohr Atom as a model for the periodic law:
En = -(Z2/n2)Ry
rn = (n2/Z)a0 =
=
energy of electrons
radius of a Bohr orbit
Structure of atoms:
Quantum numbers of electrons
Electron configurations
Core electrons and valence electrons
Periodic properties of atoms:
Energy required to remove and add an electron
Size of atoms
The atomic electron configurations of first five rows of the periodic table
Row
Configuration
First:
Second:
Third:
Fourth:
Fifth:
1s2
2s22p6
3s23p6
4s23d104p6
5s24d105p6
Shorthand
2[He]
10[Ne]
18[Ar]
36[Kr]
54[Xe]
The atomic electron configurations of first five rows of the periodic table
give the elements their signature characteristics of metals and non-metals
Row
Configuration
Shorthand
First:
Second:
Third:
Fourth:
Fifth:
1s2
2s22p6
3s23p6
4s23d104p6
5s24d105p6
2[He]
10[Ne]
18[Ar]
36[Kr]
54[Xe]
The Periodic Table built up by electron configurations: the ground
state electron configurations of the valence electrons of the elements
Shells, subshells and orbitals
Shell: a collection of orbitals with the same value of n
Example: the orbitals 3s, 3p, 3d comprise a shell with n =
3
Subshell: a collection of orbitals with the same value of
n and l
Example: for the n = 3 shell there are three subshells,
the 3s subshell, the 3p subshell and the 3d subshell
Orbital: the individual components of a shell or subshell
Example: the px, py and pz orbitals are the components of
any p (l = 1) subshell
More about Shells, subshell and orbitals
Each shell of principal quantum number n contains n subshells
n = 1, only one subshell (s)
n = 2, two subshells (s, p)
n = 3, three subshells (s, p, d)
Each subshell of quantum number l contains (2l + 1) orbitals
l = 0, (2x0 + 1) = 1 orbital (s)
l = 1, (2x1 + 1) = 3 orbitals (px, py, pZ)
l = 2, (2x2 + 1) = 5 orbitals (dxy, dyz, dxz, dx2 - y2, dz2)
The number of orbitals for a given n is n2 (solutions to wave equation)
For n = 1, one orbital; for n = 2, four orbitals, for n = 3, nine orbitals
The number of electrons that can fill a given shell = 2n2
For each orbital, there can be a maximum of 2 occupying electrons
Building up the Periodic Table
For the second row of the periodic table, the valence electrons are
electrons in the s and p orbitals: valence electrons = snpm (n less
than or equal to 2 and m less than or equal to 8)
Atom
3Li
4Be
5B
6C
7N
8O
9F
10Ne
Configuration
Core/Valence electrons
[Ne]2s ()
[Ne] 2s2 ()
[Ne] 2s22p1 ()
[Ne] 2s22p2 ()
[Ne] 2s22p3 ()
[Ne] 2s22p4 ()
[Ne] 2s22p5 ()
[Ne] 2s22p6 ()
Magnetic Properties
Paramagnetic
Closed shell (diamagnetic)
Paramagnetic
Paramagnetic
Paramagnetic
Paramagnetic
Paramagnetic
Closed shell (diamagnetic)
How do electronic configurations connect with valence electrons and
Lewis structures?
Correlation of valence electron and Lewis structures: 2p
indicates an unpaired electron in a 2p orbital
N
2
[He]2s22px2py2pz
O
2
[He]2s22px22py2pz
F
2
Ne
2
[He]2s22px22py22pz
[He]2s22px22py22pz2
Filled shell
Building up the third row of the periodic table:
From Na to Ar
Atom
11Na
12Mg
13Al
14Si
15P
16S
17Cl
18Ar
Configuration
Magnetic properties
Core/Valence electrons
[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)
The fourth row of the periodic table
Atom
Configuration
19K
18[Ar]4s
20Ca
18[Ar]4s2
_________________________________
31Ga
18[Ar]
33As
18[Ar]
32Ge
34Se
35Cl
36Kr
18[Ar]
18[Ar]
18[Ar]
18[Ar]
ten d orbitals fill up
3d104s24p1
3d104s24p2
3d104s24p3
3d104s24p4
3d104s24p5
3d104s24p6
What about
21M
through
30M?
The n + l rule: The ordering of the energies of the orbitals in a
multielectron atom increase with the value of n + l. When two
orbitals have the same value of n + l the orbital with the lower value
of n has the lower energy state.
Orbital
(n + l)
1s
2s
2p
3s
3p
4s
3d
4p
5s
(1 + 0 = 1)
(2 + 0 = 2)
(2 + 1 = 3)
(3 + 0 = 3)
(3 + 1 = 4)
(4 + 0 = 4)
(3 + 2 = 5)
(4 + 1 = 5)
(5 + 1 = 6)
Comment
Lower n (2p versus 3s) has lower energy
Lower n (3p versus 4s) has lower energy
Lower n (3d versus 4p) has lower energy
Order of filling for first four rows: 1s2 2s22p6 3s23p4 4s23d104p6 5s2
Following the (n + l) rule, the electron configurations of the transition
elements of the fourth row
21Sc
22Ti
23V
24Cr
25Mn
26Fe
27Co
28Ni
29Cu
30Zn
18[Ar]4s23d
18[Ar]4s23d2
18[Ar]4s23d3
18[Ar]4s23d4
instead 18[Ar]4s1()3d5 ()
18[Ar]4s23d5
half filled half filled
18[Ar]4s23d6
18[Ar]4s23d7
18[Ar]4s23d8
18[Ar]4s23d9 instead 18[Ar]4s1 ()3d10 ()
18[Ar]4s23d10
half filled
filled
The “surprises” for electron configurations at 24Cr and 29Cu are due to
the special stability of half filled subshells and filled subshells.
The screening of outer
electrons by the core
electrons is the basis of Zeff
The Ar atom has shells as
shown in the profile of
electron density as a function
of distance from the nucleus
This screening is the meaning
of the symbol 18[Ar]
The 4s electron of K = [Ar]4s.
is screened by the 18[Ar] core
electrons
The spatial distribution of ns
electrons as a function of r
The spatial distribution of s, p and
d electrons as a function of r
Effective nuclear charge
Effective nuclear charge, Zeff: the net positive charge attracting an electron.
An approximation to this net charge is
Zeff(effective nuclear charge) = Z(actual nuclear charge) - Zcore(core electrons)
The core electrons are in subshell between the electron in question and the nucleus. The
core electrons are said to “shield” the outer electrons from the full force of the nucleus.
Example: A 3s electron is in an orbital that is closer to the nucleus than a 3 p orbital.
Therefore, an electron in a 3s orbital is less shielded from the nucleus than the 3 p
orbital.
Rule: In many electron atoms, for a given value of n, Zeff decreases with increasing l,
because screening decreases with increasing l
For a given n: s < p < d < f
Since the energy of an orbital depends on Zeff, in a many electron atom, for a given value of
n, the energy of a orbital increases with increasing value of l.
Using electron configurations to explain Ionization Energies
Removal of an electron from a neutral atom: IE1
En = -(Z2/n2)
En = -(Z2) if n is fixed
(across a row)
En = -(n2) if Z is fixed
(down a row)
IE1
3Li
4Be
5B
6C
7N
8O
9F
10Ne
[Ne]2s ()
[Ne] 2s2 ()
[Ne] 2s22p1 ()
[Ne] 2s22p2 ()
[Ne] 2s22p3 ()
[Ne] 2s22p4 ()
[Ne] 2s22p5 ()
[Ne] 2s22p6 ()
Using electron configurations to explain Ionization Energies
Removal of the second electron from the cation: IE2
IE2 (removal from E+)
3Li
4Be
5B
6C
7N
8O
9F
10Ne
[Ne]
[Ne]
[Ne]
[Ne]
[Ne]
[Ne]
[Ne]
[Ne]
2s2 ()
2s22p1
2s22p2 ()
2s22p3 ()
2s22p4 ()
2s22p5 ()
2s22p6 ()
From the Bohr atom to all atoms: a model for the size of atoms.
r = a0(n2/Z) so that for the same value of n
r a a0(1/Zeff)
When electrons are added to the same shell (same value of n) they are
about the same distance from the nucleus as the other electrons in the
shell. The electrons in a shell with the same n are spread out and do
not shield each other from the positive charge of the nucleus very well.
Thus, the effective nuclear charge, Zeff, increases as Z increases
across the periodic table. The increasing value of Zeff draws the
electrons in closer to the nucleus, and the atom becomes more
compact.
Conclusion: The atomic radius of an atom decreases as one goes across
a period for atoms of the same value of n.
This conclusion is apparent from the formula for the radius of a Bohr
orbit:
Atomic radius: r = (n2/Z)a0
r = kn2 (r increases)if Z is fixed (going down a column)
r = k’/Z (r decreases) if n is fixed (going across a row)
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
En = -(Z2/n2)
Across row Z increases for
similar n (valence electrons
see more + as Z increases)
Down column Zeff is similar
for increasing n and r
(valence electrons further
away with same Zeff)