Table of Contents
Chapter 8 – Electron Configurations and
Chemical Periodic Trends
I.
ELECTRON SPIN ................................................................................................................................................ 2
THE 4TH QUANTUM NUMBER (MS)................................................................................................................................. 2
PARAMAGNETIC / DIAMAGNETIC BEHAVIOR ............................................................................................................... 2
II.
ELECTRON CONFIGURATIONS AND ENERGY DIAGRAMS ............................................................. 3
ENERGY LEVEL DIAGRAMS ......................................................................................................................................... 3
Energy Levels ......................................................................................................................................................... 3
3 PRINCIPLES FOR ORDER OF FILLING ENERGY LEVELS .............................................................................................. 4
ATOMIC ELECTRON CONFIGURATIONS ........................................................................................................................ 4
ELECTRON CONFIGURATIONS – EXCEPTIONS TO THE RULES ....................................................................................... 5
ION ELECTRON CONFIGURATIONS ............................................................................................................................... 6
Determining Ion Electron Configurations .............................................................................................................. 6
Predicting Oxidation States .................................................................................................................................... 6
III.
CHEMICAL PERIODIC TRENDS ................................................................................................................ 7
EFFECTIVE NUCLEAR CHARGE - ZEFF OR Z* ................................................................................................................. 8
Effective vs Inefective screening ............................................................................................................................. 9
ATOMIC SIZE ............................................................................................................................................................... 9
Trends in Atomic Size ........................................................................................................................................... 10
IONIC RADII ............................................................................................................................................................... 10
Cations .................................................................................................................................................................. 10
Anions ................................................................................................................................................................... 11
Isoelectronic Species ............................................................................................................................................ 11
IONIZATION ENERGY ................................................................................................................................................. 12
1st Ionization Energy Ei1 ...................................................................................................................................... 12
2nd Ionization Energy Ei2 ..................................................................................................................................... 12
Trends in Ionization Energy ................................................................................................................................. 13
Anomolies in Ionization Energy ............................................................................................................................ 13
ELECTRON AFFINITY ................................................................................................................................................. 14
Page 8-1
Chapter 8 – Electron Configurations and
Chemical Periodic Trends
I. Electron Spin
The 4th Quantum number (ms)
- Did not come directly from the wave equation.
- It determines the
- It can have values of
on the electron.
or
.
Pictue from McMurry Fay
- - Spinning electrons (charged species) generate a
.
- - When placed in a magnetic field they will align themselves either
the field. (They will be either
or
by or
to
the field.)
- When the first e- is assigned to an orbital, it can have either value of
.
- - When a second electron is assigned to the same orbital, it will have the
spin.
Paramagnetic / Diamagnetic Behavior
When placed into a magnetic field:
- - Diamagnetic substances are
all
the field and have
.
- - Paramagnetic substances are
at least 1
the field and have
.
- - Ferromagnetic substances are
all unpaired electrons
the field and have
with the field.
(e.g. Fe; These are permanent magnets.)
Page 8-2
II. Electron Configurations and Energy Diagrams
Energy Level Diagrams
- Energy levels increase in the same order
as the order of the
.
- For all atoms except hydrogen, the
energy associated with an orbital is
dependent on the values of
(or the sublevel).
,
Pictue from McMurry Fay
Energy Levels
Since each set of quantum numbers is unique, you would expect that the P.E. of each
orbital would also be unique, however:
- The P.E. for some orbitals are the
same and result in ____________
orbitals.
- For the H atom, all orbitals with
the same value of
are
degenerate.
- For other atoms, all orbitals with
the same values of
are
degenerate.
Sample Full Energy Level Diagram
____
____
5p
____
Pictue from McMurry Fay
____
____
____
4d
____
____
____
____
____
3d
____
____
____
5s
____
____
4p
____
____
4s
____
____
3p
____
____
2p
____
Energy spacing is not equivalent,
see above diagram.
____
3s
____
____
2s
____
1s
Page 8-3
An orbital box diagram simply lists the above energy levels sequentially:
____ ____ ____ ____ ____ ____ ____ ____ ____ ____ etc.
1s
2s
2p
3s
3p
4s
3 Principles for Order of Filling Energy Levels
Aufbau Principle
Hund’s Rule
Always fill the
energy levels first.
If you have orbitals of the same energy (
you fill 1 in each orbital before you double up. All
unpaired e- all have the
spin.
Pauli Exclusion Principle No two electrons can have the same
),
.
Atomic Electron Configurations
A summary of the orbital location of each electron is written as, for example:
Mg 1s2 2s2 2p6 3s2 etc.
The primary
QN (n) or
level
The sublevel
or orbital type
The number of
electrons in the
sublevel
Class Practice
Determine the electron configuration and orbital box diagram for:
H
He
Li
C
V
OR
OR
1 s1
1s2
1s2 2s1
1s2 2s2 2p2
1s2 2s2 2p6 3s2 3p6 4s2 3d3
[Ar] 4s2 3d3
[Ar] 3d3 4s2
This can be called
often re-written with the
electrons last
notation or a
Have class try: W
________________________________________
Page 8-4
and
notation.
An alternate trick that can be used to determine the order or sublevel filling is:
http://www.chemistry.ohio-state.edu/~grandinetti/teaching/Chem121/lectures/
Electron Configurations – Exceptions to the Rules
- There are 19 out of 109 elements whose actual electronic configuration does not
match what would be predicted by the rules that we have learned.
- This occurs because orbital energies start to get closer and closer together as we get
farther from the nucleus, and other contributions to the
of the
atom can change the filling order.
-
and
subshells are usually more stable (lower energy).
- You are responsible to know 5 of the 19 exceptions, which fall into 1 of 2
situations.
Cr and Mo
Cr [Ar] 4s2 3d4 predicted
[Ar]

4s




3d
___

4s




3d
.
becomes
[Ar]
When you move one
electron up from the 4s to
the 3d sublevel, they both
become half full .
Cr [Ar] 4s1 3d5 actual
Mo [Kr] 5s2 4d4 predicted
__________________
Page 8-5
actual
Cu, Ag, an Au
Cu [Ar] 4s2 3d9 predicted
[Ar]

4s



3d

.

4s



3d


When you move one
electron up from the 4s to
the 3d sublevel, you get one
full and one half full
sublevel.
becomes
[Ar]
Cu [Ar] 4s1 3d10 actual
Ag [Kr] 5s2 4d9 predicted
__________________ actual
Au [Xe] 6s2 5d9 predicted
__________________ actual
Ion Electron Configurations
Determining Ion Electron Configurations
Calcium ion
Ca
[Ar] 4s2

Ca2+
[Ar]
+
+
2e2e-
nitride ion
N
+ 3e2
3
[He] 2s 2p + 3e-

N3[He] 2s2 2p6
or [Ne]
species (having the same # of
electron configuration
) have the same
example: N3-, O2-, F-, Ne, Na+, Mg2+, Al3+ are isoelectronic
They have the same # of electrons, but have different numbers of protons.
All have the configuration:
.
Predicting Oxidation States
The most common oxidation states of metals is based on a gain or loss of electrons
that will leave a
electron configuration.
Iron forms two common oxidation states Fe2+ and Fe3+.
Page 8-6
Why?
Basic Rules:
Transition Metals will loose their outer
inner
electrons.
electrons before their
e.g., Ag
Ag [Kr] 5s1 4d10  Ag [Kr] 4d10 5s1  _____________________
e.g., Iron
Fe [Ar] 4s2 3d6  Fe [Ar] 3d6 4s2  ____________________
The d subshell would be more stable if half filled so Fe continues on:
Fe2+ [Ar] 3d6  _____________________
Main Group Metals loose their
electrons before their
electrons.
e.g., Lead
Pb [Xe] 6s2 4f14 5d10 6p2  Pb [Xe] 4f14 5d10 6s2 6p2  ________________________
The configuration would also be stable without the s subshell so Pb continues on:
Pb2+ [Xe] 4f14 5d10 6s2
 _________________________
III. Chemical Periodic Trends
- Many chemical and physical properties of the elements exhibit regular trends.
- Demitri Mendeleev used thes similarities to help arrange his periodic table.
Groups or families have similar behavior.
- We will try to use what we now know about electronic structure to explain
this.
All Group 1 have ______ config
“
2 have ______
“
“
3 have ______
“
similar configurations lead to
“
17 have ______
“
“
18 have ______
“
Page 8-7
similar properties.
Effective Nuclear Charge - Zeff or Z*
- The predicted nuclear charge,
, is based on the number of ___________
in the nucleus. Each positive charge adds to Z.
- Inner electrons shield the positive nucleus from the outer electrons so that
they feel less of an attractive force from the positive nucleus.
- The actual positive charge that an electron experiences is the
, Zeff or Z*
The effective charge is reduced from the full charge due to the shielding of the
nuclear charge by other electron in the atom.
This picture is taken from
http://www.chemistry.ohiostate.edu/~grandinetti/teachin
g/Chem121/lectures/
The effective nuclear charge equates the number of protons in the nucleus, Z, minus
the average number of electrons, S, between the nucleus and the electron of interest.
picture was taken from:
http://www.chemistry.ohiostate.edu/~grandinetti/
teaching/Chem121/lectures/
This
Page 8-8
Effective vs Inefective screening
- Inner shell electrons provide
(but not complete) screening,
because they lie mostly between the outer electrons and the nucleus.
- Same shell electrons provide
(negligible) screening, because
they are at a similar distance from the nucleus.
- Z* is somewhere between the true charge from the nucleus and the chage if
screening were complete.
Example:
Lithium has 3 protons, so the maximum pull from the nucleus would be
Lithium also has 3 electrons:
inner 1s electrons and
.
valence (outer) 2s e-.
If the inner electrons provided no screening, Z would be
.
If each of the inner e- could completely screen 1 proton, Z would be
.
The true Z* for Li = 1.3 (Between the two possible values.)
The concept of Effective Nuclear Charge will help us explain many of the trends
that we see across the periodic table.
Atomic Size
The boundaries of an atom are fuzzy because the locations of the electrons are
functions.
Radius is assumed to be ½ the distance between the nuclei of two covalently
bonded identical atoms.
Picture from McMurry Fay
Page 8-9
Trends in Atomic Size
a) Size gets
as you go __________
a family.
(You are adding a new set of energy levels
(shell) for each row you go down.
b) Size gets
across a period.
as you go __________
(Not what you would first expect; You are adding more electrons which repel
each other so you would expect the radius to
.
This is overshadowed by the increase in
as you go across the table.
(You are adding to the
energy level so shielding from the new electrons
is
. However, the increased number of
(+) charges
in the nucleus attracts the electrons closer in.
When you get to the end of a period or row, the size
dramatically
because you are now adding to a new
and ALL of the
previous electrons now provide
shielding.
Ionic Radii
Cations
Cations always have a
radius that the parent atom.
- The same number of (+) charges are now distributed between a ____________
number of (-) charges. (The electrons are held
closely.)
- There are
e- to e- repulsive forces.
- The larges change occurs when the e- removed is the last electron in the outer
.
Ca > Ca+
Ca >>> Ca 2+
[Ar] 4s2  [Ar] 4s1
[Ar] 4s2  [Ar]
(The 4th energy level is no
longer used.)
Page 8-10
Example:
The radius of Na = 154 pm; Na+ = 116 pm
Picture from McMurry Fay
Anions
Anions always have a
radius that the parent atom.
- The same number of (+) charges are now distributed between a ____________
number of (-) charges. (The electrons are held
closely.)
- There are
e- to e- repulsive forces.
Example:
The radius of Cl = 97 pm;
The radius of Cl- = 167 pm
Picture from McMurry Fay
Isoelectronic Species
Isoelectronic species are ones that have the same number of
different numbers of
.
As the number of protons goes
goes
.
, the attractive force goes
Page 8-11
but
and the radius
Ionization Energy
The ionization energy is the energy required to remove
atom or ion in the
state.
electron from an
1st Ionization Energy Ei1
E is always
X(g)  X+(g) + 1 e-
.
2nd Ionization Energy Ei2
E is always
X+ (g)  X2+(g) + 1 e-
than Ei1.
3rd, 4th, 5th etc ionization energies can also be calculated.
Example:
1st
2nd
3rd
4th
Al(g)  Al+(g) + 1 eAl+ (g)  Al2+(g) + 1 eAl2+ (g)  Al3+(g) + 1 eAl3+ (g)  Al4+(g) + 1 e-
I.E. (kJ/mol)
580
1,815
2,740
116,000
Electron Config.
[Ne] 3s2 3p1  [Ne] 3s2
[Ne] 3s2  [Ne] 3s1
[Ne] 3s1  [Ne]
[Ne]  [He] 2s2 2p5
Why is the 4th I.E. so large?
You are trying to remove an electron from a noble gas electron configuration
which is very
.
Page 8-12
Trends in Ionization Energy
a) I.E.
as you go ________
across the periodic table.
Zeff goes up as you go right (more
positive charge and ineffective
shielding because each new e- is in
the same shell. It gets harder and
harder to remove an electron that
is held more tightly.
b) I.E.
as you go
the group or family.
You are removing electrons that are farther out in shells that are farther from
the nucleus and are better screened from the pull of the nucleus.
Anomolies in Ionization Energy
Picture from McMurry Fay
Unusually high I.E.s are observed for Be, N, Mg, and P.
These high I.E.s for some elements can be explained as follows:
Be/Mg: You are removing an e- from a filled (stable)
orbital.
N/P: You are removing an e- from a ½ filled (stable)
sublevel.
Elements with low ionization energies make good
(They give up an electron relatively easily.)
Page 8-13
agents.
Electron Affinity
Electron afinity is defined as the energy change that occurs when an electron is
to the lowest energy unoccupied orbital in a
atom.
X(g) + 1 e- X-(g)
Adding a second electron
E can be _________________________
E is always
due to repulsive forces.
Examples:
F(g) + e-  F-(g)
E.A. –328 kJ/mol
N(g) + e-  N-(g)
E.A. +7 kJ/mol
Ca(g) + e-  Ca-(g)
E.A. –2 kJ/mol
Ne(g) + e-  Ne-(g)
E.A. +116 kJ/mol
F becomes more stable (attains a noble gas configuration).
N & Ca stabilities are not significantly changed.
Ne becomes less stable (looses the noble gas configuration)
Note: Our text assigns a negative energy to an EA that is favored and a
positive energy to one that is not spontaneous to be consistent with the other
energy related terms we have studied. Many other textsreverse the signs so
that an atom that wants an electron has a high positive value. We need to
realize that when we say that an element has a
we are talking about it having a large
Page 8-14
electronegativity,
value!