Transcript Chapter 8 Electron Configuration and Chemical Periodicity 8-1
8-1
Chapter 8 Electron Configuration and Chemical Periodicity
8-2
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
Table 8.1
Mendeleev’s
Predicted
Properties of Germanium (“eka Silicon”) and Its Actual Properties
8-3 Property atomic mass appearance density molar volume specific heat capacity oxide formula oxide density sulfide formula and solubility chloride formula (boiling point) chloride density element preparation Predicted Properties of eka Silicon(E) 72 amu gray metal 5.5 g/cm
3
13 cm
3
/mol 0.31 J/g
.
K EO
2
4.7 g/cm
3
ES
2
; insoluble in H
2
O; soluble in aqueous (NH
4
)
2
S ECl
4
; (< 100
o
C) 1.9 g/cm
3
reduction of K
2
EF
6
sodium with Actual Properties of Germanium (Ge) 72.61 amu gray metal 5.32 g/cm
3
13.65 cm
3
/mol 0.32 J/g
.
K GeO
2
4.23 g/cm
3
GeS
2
; insoluble in H
2
O; soluble in aqueous (NH
4
)
2
S GeCl
4
; (84
o
C) 1.844 g/cm sodium
3
reduction of K
2
GeF
6
with
Observing the Effect of Electron Spin
8-4 Figure 8.1
8-5
Table 8.2 Summary of Quantum Numbers of Electrons in Atoms
Name
principal angular momentum magnetic
Symbol
n
l
m
l
Allowed Values
positive integers (1, 2, 3,…) integers from 0 to
n
-1
Property
orbital energy (size) integers from -
l
to 0 to +
l
orbital shape (
l
values of 0, 1, 2 and 3 correspond to
s
,
p
,
d
and
f
orbitals, respectively.) orbital orientation spin
m
s
+1/2 or -1/2 direction of e
-
spin Each electron in an atom has its own unique set of four (4) quantum numbers.
8-6
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)
Spectral evidence of energy-level splitting in many-electron systems
8-7 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 on its
l
value (shape)
n
value (size) and somewhat
Factors Affecting Atomic Orbital Energies
Effect of nuclear charge (Z
effective
)
Higher nuclear charge lowers orbital energy (stabilizes the system) by increasing nucleus-electron attractions.
8-8
Effect of electron repulsions (shielding)
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.
8-9
The effect of nuclear charge
Greater nuclear charge lowers orbital energy Figure 8.3
The effect of another electron in the same orbital
Each electron shields the other from the full nuclear charge, thus raising orbital energy 8-10 Figure 8.4
QuickTime™ and a Photo - JPEG decompressor are needed to see this picture.
The effect of other electrons in inner orbitals
Inner electrons shield outer electrons very well and raise orbital energy greatly 8-11 Figure 8.5
QuickTime™ and a Photo - JPEG decompressor are needed to see this picture.
Figure 8.6
The effect of orbital shape
2
s
electron farther from nucleus than 2
p
electron, but penetrates near nucleus; increased attraction results in lower orbital energy 8-12
8-13 General Rule for Predicting Relative Sublevel Energies For a given
n
value, the lower the
l
energy; thus….
value, the lower the sublevel
s < p < d < f
8-14 Figure 8.7
Order for filling energy sublevels with electrons
Illustrating Orbital Occupancies
A. The electron configuration
n
l
# of electrons in the sublevel as
s
,
p
,
d
or
f
B. The orbital diagram (box or circle)
A vertical orbital diagram for the Li ground state
Sublevel energy increases from bottom to top
1
s
2
2
s
1
8-15 Figure 8.8
light = half-filled no color = empty dark = filled, spin-paired
8-16
Hund’s Rule
When orbitals of equal energy are available, the electron configuration of lowest energy has the maximum number of
unpaired
electrons with parallel spins.
Sample Problem 8.1
Determining Quantum Numbers from Orbital Diagrams PROBLEM:
Write a set of quantum numbers for the third electron and a set for the eighth electron of the fluorine (F) atom.
PLAN:
Use the orbital diagram to find the third and eighth electrons.
9 F
Up arrow = +1/2 Down arrow = -1/2
1
s
2
s
2
p
SOLUTION:
The third electron is in the 2
s
orbital. Its quantum numbers are:
n
= 2
l
= 0
m
l
= 0
m
s
= +1/2 The eighth electron is in a 2
p
orbital. Its quantum numbers are:
n
= 2
l
= 1
m
l
= -1
m
s
= -1/2 8-17
Orbital occupancy for the first 10 elements, H through Ne
8-18 Figure 8.9
He and Ne have filled outer shells: confers chemical inertness
8-19
Condensed ground-state electron configurations in the first three periods
8-20 Figure 8.10
Similar outer electron configurations correlate with similar chemical behavior.
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
recur periodically.
that QuickTime™ and a Photo - JPEG decompressor are needed to see this picture.
Figure 8.11
8-21
8-22 Cr and Cu: Half-filled and filled sublevels are unexpectedly stable!
8-23
Figure 8.12
A periodic table of partial ground-state electron configurations
8-24
8-25
The relation between orbital filling and the Periodic Table
Figure 8.13
Common tool used to predict the filling order Of sublevels
n
values are constant horizontally
l
values are constant vertically combined values of
n
+1 are constant diagonally p. 302 8-26
8-27
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
8-28
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 (
n
2
) gives the total number of orbitals in that energy level; 2
n
2
gives the maximum number of electrons in the energy level.
SAMPLE PROBLEM 8.2
Determining Electron Configuration PROBLEM:
Using the periodic table, give the full and condensed electron configurations, partial orbital diagrams showing valence electrons, and number of inner electrons for the following elements:
(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) full configuration: condensed configuration: 1
s
2
2
s
2
2 [Ar] 4
s
1
p
6
3
s
2
3
p
6
4
s
1
K has 18 inner electrons.
partial orbital diagram: 4
s
1
8-29
SAMPLE PROBLEM 8.2:
(continued) 8-30
(b)
for Mo (
Z
= 42) full configuration: condensed configuration: partial orbital diagram: 1
s
2
2
s
2
2
p
6
3
s
2
3
p
6
4
s
2
3
d
10
4
p
6
5
s
1
4
d
5
[Kr] 5
s
1
4
d
5
Mo has 36 inner electrons and 6 valence electrons.
5
s
1
4
d
5 (c)
for Pb (
Z
= 82) full configuration: 1
s
2
2
s
2
2
p
6
3
s
2
3
p
6
4
s
2
3
d
10
4
p
6
5
s
2
4
d
10
5
p
6
6
s
2
condensed configuration: [Xe] 6
s
2
4
f
14
5
d
10
6
p
2
partial orbital diagram: Pb has 78 inner electrons and 4 valence electrons.
4
f
14
5
d
10
6
p
2
6
s
2
6
p
2
8-31
KEY PRINCIPLE
All physical and chemical properties of the elements are based on the electronic configurations of their atoms.
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
8-32
Atomic radii of main-group and transition elements Opposing forces
: Changes in
n
and changes in
Z
eff Overall Trends
(A)
n
dominates within a group; atomic radius generally
increases
in a group from top to bottom (B)
Z
eff
dominates within a period; atomic radius generally
decreases
in a period from left to right Figure 8.15
8-33
Large size shifts when moving from one period to the next
Periodicity of atomic radius
8-34 Figure 8.16
SAMPLE PROBLEM 8.3
Ranking Elements by Atomic Size PROBLEM:
Using only the periodic table, rank each set of main group elements in order of decreasing 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 > Ca > Mg
(b)
K > Ca > Ga
(c)
Rb > Br > Kr
(d)
Rb > Sr > Ca These elements are in Group 2A.
These elements are in Period 4.
Rb has a higher energy level and is far to the left. Br is to the left of Kr.
Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.
8-35
8-36
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 IE
1
= first ionization energy : removes an outermost electron from the gaseous atom: atom(
g
) ion
+
(
g
) + e
-
∆E = IE
1
> 0 IE
2
= second ionization energy : removes a second electron from the gaseous ion: ion
+
(
g
) ion
+ 2
(
g
) + e
-
∆E = IE
2
> IE
1
Atoms with a low IE
1
with a high IE
1
tend to form cations during reactions, whereas those (except noble gases) often form anions.
8-37 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.
Periodicity of first ionization energy (IE 1 )
Lowest values for alkali metals; highest values for noble gases 8-38 Figure 8.17
First ionization energies of the main-group elements
Increase within a period and decrease within a group Figure 8.18
8-39
SAMPLE PROBLEM 8.4
Ranking Elements by First Ionization Energy PROBLEM: (a)
Kr, He, Ar Using the periodic table, rank the elements in each of the following sets in order of decreasing IE
1
:
(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 > Ar > Kr Group 8A elements- IE decreases down a group.
8-40
(b)
Te > Sb > Sn Period 5 elements - IE increases across a period.
(c)
Ca > K > Rb Ca is to the right of K; Rb is below K.
(d)
Xe > I > Cs I is to the left of Xe; Cs is further to the left and down one period.
8-41 Figure 8.19
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, IE
3
)
8-42
SAMPLE PROBLEM 8.5
Identifying an Element from Successive Ionization Energies PROBLEM:
Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration: IE
1
1012 IE
2
1903 IE
3
2910 IE
4
4956 IE
5
6278 IE
6
22,230
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 IE
6
, that is, after the 5th valence electron has been removed. The element must have five valence electrons with a valence configuration of 3s
2
3p
3
, The element must be phosphorus. P (
Z
= 15).
The complete electronic configuration is: 1
s
2
2
s
2
2
p
6
3
s
2
3
p
3
.
8-43
8-44
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 = EA
1
( usually negative ) EA
2
is always positive (adding negative charge to negatively charged ion).
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 8-45 Figure 8.20
8-46
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)
8-47
Trends in three atomic properties
Figure 8.21
8-48
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
Figure 8.22
8-49
Trends in metallic behavior
8-50 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
8-51
The change in metallic behavior in Group 5A(15) and Period 3
Figure 8.23
8-52
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).
The trend in acid-base behavior of element oxides
Figure 8.24
red = oxides are acidic blue = oxides are basic 8-53
8-54
Examples
Na
2
O(
s
) + H
2
O(
l
) 2NaOH(
aq
) N
2
O
5
(
s
) + H
2
O(
l
) 2HNO
3
(
aq
) P
4
O
10
(
s
) + 6H
2
O(
l
) 4H
3
PO
4
(
aq
) Bi
2
O
3
(
s
) + 6HNO
3
(
aq
) 2Bi(NO
3
)
3
(
aq
) + 3H
2
O(
l
) Amphoteric Behavior Al
2
O
3
(s) + 6HCl(
aq
) 2AlCl
3
(
aq
) + 3H
2
O(
l
) Al
2
O
3
(
s
) + 2NaOH(
aq
) + 3H
2
O(
l
) 2NaAl(OH)
4
(
aq
)
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 .
Sn ([Kr]5
s
2
4
d
10
5
p
2
) Sn ([Kr]5
s
2
4
d
10
5
p
2
) Sn
+4
([Kr]4
d
10
) + 4e
-
Sn
+2
([Kr]5
s
2
4
d
10
) + 2e
-
8-55
Main-group ions and the noble gas configurations
8-56 Figure 8.25
SAMPLE PROBLEM 8.6
Writing Electron Configurations of Main-Group Ions PROBLEM:
Using condensed electron configurations, write reactions for the formation of the common ions of the following elements:
(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: (a)
Iodine (
Z
= 53) is in Group 7A and will gain one e
-
I([Kr]5
s
2
4
d
10
5
p
5
) + e
-
I
-
to be isoelectronic with Xe: ([Kr]5
s
2
4
d
10
5
p
6
)
(b)
Potassium (
Z
= 19) is in Group 1A and will lose one e
-
K ([Ar]4
s
1
) K
+
([Ar]) + e
-
to be isoelectronic with Ar:
(c)
Indium (
Z
= 49) is in Group 3A(13) and can lose either one electron or three electrons: In ([Kr]5
s
2
4
d
10
5
p
1
) In
+
([Kr]5
s
2
4
d
10
) + e
-
In ([Kr]5
s
2
4
d
10
5
p
1
) In
3+
([Kr]4
d
10
) + 3e
-
8-57
8-58
Monatomic Ions
Transition Metal Ions
Rarely attain a noble gas configuration Form more than one cation by losing all of their
ns
some of their (
n
-1)
d
electrons and For Period 4 transition metals, the 4
s
stable that the 3
d
orbital is more orbitals; thus the rule “first in, first out” applies.
The Period 4 crossover in sublevel energies
8-59 Figure 8.26
8-60
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 (
n
-1)
d ns
electrons before electrons Non-metals : add electrons to the
p
orbital of highest
n
value
8-61
Magnetic Properties of Transition Metal Ions
Chemical species (atoms, ions, molecules) with one or more unpaired electrons are affected by external magnetic fields.
Ag (
Z
=47) [Kr]5
s
1
4
d
10
Cd (
Z
=48) [Kr]5
s
2
4
d
10
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).
Apparatus for measuring the magnetic behavior of a sample
8-62 Figure 8.27
8-63
Some Examples
Fe
+3
exhibits greater paramagnetism than Fe.
Fe ([Ar]4
s
2
3
d
6
) Fe
+3
([Ar]3
d
5
) + 3e
-
Zn, Zn
+2
and Cu
+
are diamagnetic, but Cu is paramagnetic.
Cu ([Ar]4
s
1
3
d
10
) Cu
+
([Ar]3
d
10
) + e
-
Zn ([Ar]4
s
2
3
d
10
) Zn
+2
([Ar]3
d
10
) + 2e
-
SAMPLE PROBLEM 8.7
Writing Electron Configurations and Predicting Magnetic Behavior of Transition Metal Ions PROBLEM:
Use condensed electron configurations, write the reaction for the formation of each transition metal ion and predict whether the ion is paramagnetic.
(a)
Mn
2+
(
Z
= 25)
(b)
Cr
3+
(
Z
= 24)
(c)
Hg
2+
(
Z
= 80)
PLAN: SOLUTION:
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.
(a)
Mn
2+
(
Z
= 25) Mn([Ar]4
s
2
3
d
5
) Mn
2+
([Ar]3
d
5
) + 2e
-
paramagnetic
(b)
Cr
3+
(
Z
= 24) Cr([Ar]4
s
1
3
d
5
) Cr
3+
([Ar] 3
d
3
) + 3e
-
paramagnetic 8-64
(c)
Hg
2+
(
Z
= 80) Hg([Xe]6
s
2
4
f
14
5
d
10
) Hg
2+
([Xe] 4
f
14
5
d
10
) + 2e
-
not paramagnetic (is diamagnetic)
8-65
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).
Anions are
larger
than their parent atoms (increase in electron-electron repulsions).
Depicting ionic radii
8-66 Figure 8.28
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 8-67 Figure 8.29
8-68
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
SAMPLE PROBLEM 8.8
Ranking Ions by Size PROBLEM: PLAN:
Rank each set of ions in order of
decreasing
size, and explain your ranking:
(a)
Ca
2+
, Sr
2+
, Mg
2+ (b)
K
+
, S
2 -
, Cl
(c)
Au
+
, Au
3+
Compare positions in the periodic table, formation of positive and negative ions and changes in size due to gain or loss of electrons.
SOLUTION: (a) (b)
Sr
2+
S
2 -
> Ca > Cl
2+
> K > Mg
+ 2+
These are members of the same Group (2A) and therefore decrease in size going up the group.
These ions are isoelectronic; S
2 -
therefore is the largest while K
+
has the smallest Z
eff
and is a cation with a large Z
eff
and is the smallest.
(c)
Au
+
> Au
3+
The higher the positive charge, the smaller the ion.
8-69