Electron Configuration and Chemical Periodicity

1
Chapter 8
Electron Configuration and Chemical Periodicity
2
Electron Configuration and Chemical Periodicity
8.1 Development of the Periodic Table
8.2 Characteristics of Many-Electron Atoms
8.3 The Quantum-Mechanical Model and the
Periodic Table
8.4 Trends in Some Key Periodic Atomic
Properties
8.5 The Connection Between Atomic Structure and
Chemical Reactivity
3
Predicted Properties of eka Silicon(E)
Actual Properties of Germanium (Ge)
Property
atomic mass
72 amu
72.61 amu
appearance
gray metal
gray metal
density
5.5 g/cm3
5.32 g/cm3
molar volume
13 cm3/mol
13.65 cm3/mol
specific heat capacity
0.31 J/g.K
0.32 J/g.K
oxide formula
EO2
GeO2
oxide density
4.7 g/cm3
4.23 g/cm3
sulfide formula and solubility
ES2 insoluble in H2O soluble in aqueous (NH4)2S
GeS2 insoluble in H2O soluble in aqueous (NH4)2S
ECl4 (lt 100 oC)
chloride formula (boiling point)
GeCl4 (84 oC)
chloride density
1.9 g/cm3
1.844 g/cm3
element preparation
reduction of K2EF6 with sodium
reduction of K2GeF6 with sodium
4
Observing the Effect of Electron Spin
Figure 8.1
5
Table 8.2 Summary of Quantum Numbers of
Electrons in Atoms
Name
Symbol
Allowed Values
Property
Each electron in an atom has its own unique set
of four (4) quantum numbers.
6
The Pauli Exclusion Principle
No two electrons in the same atom can have the
same four quantum numbers
An atomic orbital can hold a maximum of two
electrons and they must have opposite spins
(paired spins)
7
Spectral evidence of energy-level splitting in
many-electron systems
Figure 8.2
Many-electron atoms have nucleus-electron and
electron-electron interactions
Leads to the splitting of energy levels into
sublevels of differing energies the energy of
an orbital depends mostly on its n value (size)
and somewhat on its l value (shape)
8
Factors Affecting Atomic Orbital Energies
Higher nuclear charge lowers orbital energy
(stabilizes the system) by increasing
nucleus-electron attractions.
1. Additional electron in the same orbital
An additional electron raises the orbital energy
through electron-electron repulsions.
2. Additional electrons in inner orbitals
Inner electrons shield outer electrons more
effectively than do electrons in the same
sublevel.
9
The effect of nuclear charge
Greater nuclear charge lowers orbital energy
Figure 8.3
10
The effect of another electron in the same orbital
Each electron shields the other from the
full nuclear charge, thus raising orbital energy
Figure 8.4
11
The effect of other electrons in inner orbitals
Inner electrons shield outer electrons very
well and raise orbital energy greatly
Figure 8.5
12
The effect of orbital shape
Figure 8.6
2s electron farther from nucleus than 2p
electron, but penetrates near nucleus increased
attraction results in lower orbital energy
13
General Rule for Predicting Relative Sublevel
Energies
For a given n value, the lower the l value, the
lower the sublevel energy thus.
s lt p lt d lt f
14
Figure 8.7
Order for filling energy sublevels with electrons
Illustrating Orbital Occupancies
A. The electron configuration

of electrons in the sublevel
n
l
B. The orbital diagram (box or circle)
15
A vertical orbital diagram for the Li ground state
no color empty
light half-filled
Sublevel energy increases from bottom to top
1s22s1
dark filled, spin-paired
Figure 8.8
16
Hunds Rule
When orbitals of equal energy are available, the
electron configuration of lowest energy has the
maximum number of unpaired electrons with
parallel spins.
17
Determining Quantum Numbers from Orbital Diagrams
Sample Problem 8.1
PLAN
Use the orbital diagram to find the third and
eighth electrons.
Up arrow 1/2 Down arrow -1/2
SOLUTION
The third electron is in the 2s orbital. Its
quantum numbers are
2
0
0
1/2
The eighth electron is in a 2p orbital. Its
quantum numbers are
2
-1
-1/2
1
18
Orbital occupancy for the first 10 elements, H
through Ne
Figure 8.9
He and Ne have filled outer shells confers
chemical inertness
19
(No Transcript)
20
Condensed ground-state electron configurations in
the first three periods
Figure 8.10
Similar outer electron configurations correlate
with similar chemical behavior.
21
Similar Reactivities within A Group
Orbitals are filled in order of
increasing energy, which leads to outer
electron configurations that recur
periodically, which leads to chemical properties
that recur periodically.
Figure 8.11
22
Cr and Cu Half-filled and filled sublevels are
unexpectedly stable!
23
(No Transcript)
24
A periodic table of partial ground-state electron
configurations
Figure 8.12
25
The relation between orbital filling and the
Periodic Table
Figure 8.13
26
Common tool used to predict the filling order Of
sublevels
n values are constant horizontally l values are
constant vertically combined values of n1 are
constant diagonally
p. 302
27
Categories of Electrons
Inner (core) electrons fill all the lower
energy levels of an atom
Outer electrons those electrons in the highest
energy level (highest n value) of an atom
Valence electrons those involved in forming
compounds the bonding electrons among the
main-group elements, the valence electrons are
the outer electrons
28
General Observations about the Periodic Table
A. The group number equals the number of outer
electrons (those with the highest value of n)
(main-group elements only)
B. The period number is the n value of the
highest energy level.
C. The n value squared (n2) gives the total
number of orbitals in that energy level 2n2
gives the maximum number of electrons in the
energy level.
29
SAMPLE PROBLEM 8.2
Determining Electron Configuration
(a) potassium (K Z 19)
(b) molybdenum (Mo Z 42)
(c) lead (Pb Z 82)
PLAN
Use the atomic number for the number of electrons
and the periodic table for the order of filling
of the electron orbitals. Condensed
configurations consist of the preceding noble gas
plus the outer electrons.
SOLUTION
(a) for K (Z 19)
1s22s22p63s23p64s1
Ar 4s1
K has 18 inner electrons.
30
SAMPLE PROBLEM 8.2
(continued)
(b) for Mo (Z 42)
1s22s22p63s23p64s23d104p65s14d5
Kr 5s14d5
Mo has 36 inner electrons and 6 valence electrons.
(c) for Pb (Z 82)
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2
Xe 6s24f145d106p2
Pb has 78 inner electrons and 4 valence electrons.
31
KEY PRINCIPLE
All physical and chemical properties of the
elements are based on the electronic
configurations of their atoms.
32
Defining metallic and covalent radii
A. Metallic radius 1/2 the distance between
adjacent nuclei in a crystal
B. Covalent radius 1/2 the distance between
bonded nuclei in a molecule
Figure 8.14
33
Atomic radii of main-group and transition elements
Opposing forces Changes in n and changes in Zeff
Overall Trends (A) n dominates within a group
atomic radius generally increases in a group
from top to bottom (B) Zeff dominates within a
period atomic radius generally decreases in a
period from left to right
Figure 8.15
34
Periodicity of atomic radius
Large size shifts when moving from one period to
the next
Figure 8.16
35
SAMPLE PROBLEM 8.3
Ranking Elements by Atomic Size
(a) Ca, Mg, Sr
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
PLAN
Size increases down a group size decreases
across a period.
SOLUTION
(a) Sr gt Ca gt Mg
These elements are in Group 2A.
(b) K gt Ca gt Ga
These elements are in Period 4.
(c) Rb gt Br gt Kr
Rb has a higher energy level and is far to the
left. Br is to the left of Kr.
(d) Rb gt Sr gt Ca
Ca is one energy level smaller than Rb and Sr.
Rb is to the left of Sr.
36
Ionization Energy
The amount of energy required for the complete
removal of 1 mol of electrons from 1 mol of
gaseous atoms or ions an energy- requiring
process value is positive in sign
IE1 first ionization energy removes an
outermost electron from the gaseous atom
atom(g) ion(g) e- ?E IE1 gt 0

IE2 second ionization energy removes a second
electron from the gaseous ion ion(g)
ion2(g) e- ?E IE2 gt IE1
Atoms with a low IE1 tend to form cations during
reactions, whereas those with a high IE1 (except
noble gases) often form anions.
37
Ionization Energies Correlations with Atomic
Size
1. As size decreases, it take more energy to
remove an electron.
2. Ionization energy generally decreases down a
group.
3. Ionization generally increases across a
period.

38
Periodicity of first ionization energy (IE1)
Figure 8.17
Lowest values for alkali metals highest values
for noble gases
39
First ionization energies of the main-group
elements
Increase within a period and decrease within a
group
Figure 8.18
40
SAMPLE PROBLEM 8.4
Ranking Elements by First Ionization Energy
(a) Kr, He, Ar
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
PLAN
IE decreases down in a group IE increases
across a period.
SOLUTION
(a) He gt Ar gt Kr
Group 8A elements- IE decreases down a group.
(b) Te gt Sb gt Sn
Period 5 elements - IE increases across a period.
(c) Ca gt K gt Rb
Ca is to the right of K Rb is below K.
(d) Xe gt I gt Cs
I is to the left of Xe Cs is further to the
left and down one period.
41
The first three ionization energies of beryllium
(in MJ/mol)
Successive IEs increase, but a large increase is
observed to remove the first core electron (for
Be, IE3)
Figure 8.19
42
(No Transcript)
43
SAMPLE PROBLEM 8.5
Identifying an Element from Successive Ionization
Energies
PLAN
Look for a large increase in energy that
indicates that all of the valence electrons have
been removed.
SOLUTION
The largest increase occurs at IE6, that is,
after the 5th valence electron has been removed.
The element must have five valence electrons with
a valence configuration of 3s23p3, The element
must be phosphorus. P (Z 15). The complete
electronic configuration is 1s22s22p63s23p3.
44
Electron Affinity (EA)
The energy change accompanying the addition of 1
mol of electrons to 1 mol of gaseous atoms or
ions.
atom(g) e- ion-(g) ?E
EA1 (usually negative)
EA2 is always positive (adding negative charge to
negatively charged ion).
45
Electron affinities of the main-group elements
Negative values energy is released when the ion
forms
Positive values energy is absorbed to form the
anion
Figure 8.20
46
General Trends Involving IEs and EAs
Reactive non-metals Groups 6A and 7A in their
ionic compounds they form negative ions (have
high IEs and very (-) EAs)
Reactive metals Group 1A in their ionic
compounds, they form positive ions (have low
IEs and slightly (-) EAs)
Noble gases Group 8A they do not lose or gain
electrons (have very high IEs and slightly ()
EAs)
47
Trends in three atomic properties
Figure 8.21
48
Metallic Behavior
Metals shiny solids tend to lose electrons in
reactions with non-metals (left and lower 3/4 of
periodic table)
Non-metals tend to gain electrons in reactions
with metals upper right-hand quarter of
periodic table
Metalloids have intermediate properties
located between the metals and non-metals in
the periodic table
Metallic behavior decreases left to right and
increases top to bottom in the periodic table
49
Trends in metallic behavior
Figure 8.22
50
Moving down a GROUP elements at the top tend to
form anions and those at the bottom tend to form
cations
Moving across a PERIOD elements at the left
tend to form cations and those at the right lend
to form anions
51
The change in metallic behavior in Group 5A(15)
and Period 3
Figure 8.23
52
Acid-Base Properties
Main-group metals transfer electrons to oxygen
their oxides are ionic in water these oxides act
as bases (produce OH-)
Nonmetals share electrons with oxygen their
oxides are covalent in water these oxides act as
acids (produce H)
Some metals and many metalloids form oxides that
are amphoteric (can act as an acid or a base in
water).
53
The trend in acid-base behavior of element oxides
Figure 8.24
red oxides are acidic blue oxides are
basic
54
Examples
Na2O(s) H2O(l) 2NaOH(aq)
N2O5(s) H2O(l) 2HNO3(aq)
P4O10(s) 6H2O(l) 4H3PO4(aq)
Bi2O3(s) 6HNO3(aq) 2Bi(NO3)3(aq)
3H2O(l)
Amphoteric Behavior
Al2O3(s) 6HCl(aq) 2AlCl3(aq)
3H2O(l)
Al2O3(s) 2NaOH(aq) 3H2O(l)
2NaAl(OH)4(aq)
55
Monatomic Ions
Main Group
Elements in Groups 1A, 2A, 6A and 7A that readily
form ions either lose or gain electrons to attain
a filled outer level and thus a noble gas
configuration. Their ions are said to be
isoelectronic with the nearest noble gas.
Elements in Groups 3A, 4A and 5A form cations via
a different process they attain pseudo-noble gas
configurations.
Sn4 (Kr4d10) 4e-
Sn (Kr5s24d105p2)
Sn2 (Kr5s24d10) 2e-
Sn (Kr5s24d105p2)
56
Main-group ions and the noble gas configurations
Figure 8.25
57
SAMPLE PROBLEM 8.6
Writing Electron Configurations of Main-Group Ions
(a) iodine (Z 53)
(b) potassium (Z 19)
(c) indium (Z 49)
PLAN
Ions of elements in Groups 1A, 2A, 6A and 7A are
usually isoelectronic with the nearest noble gas.
Metals in Groups 3A to 5A can lose the np, or ns
and np, electrons.
SOLUTION
58
Monatomic Ions
Transition Metal Ions
Rarely attain a noble gas configuration
Form more than one cation by losing all of their
ns and some of their (n-1)d electrons
For Period 4 transition metals, the 4s orbital is
more stable that the 3d orbitals thus the rule
first in, first out applies.
59
The Period 4 crossover in sublevel energies
Figure 8.26
60
General Rules For Ion Formation
Main group, s-block metals remove all electrons
with highest n value
Main group, p-block metals remove np electrons
before ns electrons
Transition (d-block) metals remove ns electrons
before (n-1)d electrons
Non-metals add electrons to the p orbital of
highest n value
61
Magnetic Properties of Transition Metal Ions
Chemical species (atoms, ions, molecules) with
one or more unpaired electrons are affected by
external magnetic fields.
Ag (Z47) Kr5s14d10 Cd (Z48) Kr5s24d10
Species with unpaired electrons exhibit
paramagnetism (attracted by an external magnetic
field).
Species with all electrons paired exhibit
diamagnetism (not attracted by an external
magnetic field).
62
Apparatus for measuring the magnetic behavior of
a sample
Figure 8.27
63
Some Examples
Fe3 exhibits greater paramagnetism than Fe.
Fe (Ar4s23d6) Fe3 (Ar3d5) 3e-
Zn, Zn2 and Cu are diamagnetic, but Cu is
paramagnetic.
Cu (Ar4s13d10) Cu (Ar3d10)
e- Zn (Ar4s23d10) Zn2 (Ar3d10)
2e-
64
Writing Electron Configurations and Predicting
Magnetic Behavior of Transition Metal Ions
SAMPLE PROBLEM 8.7
(a) Mn2(Z 25)
(b) Cr3(Z 24)
(c) Hg2(Z 80)
PLAN
Write the electron configuration and remove
electrons starting with the ns electrons to
attain the ion charge. If the remaining
configuration has unpaired electrons, the ion is
paramagnetic.
SOLUTION
paramagnetic
paramagnetic
not paramagnetic (is diamagnetic)
65
Ionic Size vs Atomic Size
Ionic radius an estimate of the size of an ion
in a crystalline ionic compound
General Observations
Cations are smaller than their parent atoms
(decrease in electron-electron repulsions). Anion
s are larger than their parent atoms (increase in
electron-electron repulsions).
66
Depicting ionic radii
Figure 8.28
67
Ionic vs atomic radius
Ionic size increases down a group
Trends in periods are complex
For atoms that form more than one cation
the greater the ionic charge, the smaller the
ionic radius
Figure 8.29
68
Summary on Ionic Size
Ionic size increases down a group. Ionic size
decreases across a period but increases from
cation to anion. Ionic size decreases with
increasing () (or decreasing (-)) charge in an
isoelectronic series Ionic size decreases as
charge increases for different cations of
a given element
69
SAMPLE PROBLEM 8.8
Ranking Ions by Size
(a) Ca2, Sr2, Mg2
(b) K, S2-, Cl-
(c) Au, Au3
PLAN
Compare positions in the periodic table,
formation of positive and negative ions and
changes in size due to gain or loss of electrons.
SOLUTION
These are members of the same Group (2A) and
therefore decrease in size going up the group.
(a) Sr2 gt Ca2 gt Mg2
These ions are isoelectronic S2- has the
smallest Zeff and therefore is the largest while
K is a cation with a large Zeff and is the
smallest.
(b) S2- gt Cl- gt K
(c) Au gt Au3
The higher the positive charge, the smaller the
ion.