Orbital Diagrams
Element
Total
Electrons
1s
H
1
He
2
Li
3
Be
4
2s
2p
3s
Orbital Diagrams
Element
Total
Electrons
1s
B
5
C
6
N
7
O
8
F
9
2s
2p
3s
Orbital Diagrams
For Ne (10 e-)
1s
2s
2p
3s
Filling of the 2p subshell is complete at neon.
The outermost shell (n = 2) contains an
octet (8) of electrons.
Orbital Diagrams
 Every noble gas has a complete outer shell.
He:
2 electrons in the outer shell
 1s2
All other noble gases : Ne, Ar, Kr, Xe, Rn
 an octet of electrons in the outer shell
 s2 p6
Orbital Diagrams
 This configuration is exceptionally stable.
Responsible for the unreactive nature of
the noble gases.
Elements that ionize easily do so in a way
that gives them the same octet of
electrons: They become “isoelectric” with
a noble gas!
Orbital Diagrams
For Sodium (Na)
11 electrons
1 more electron than the noble gas
neon
Valence e1s
2s
2p
Neon core
3s
highest
energy
level e-
Orbital Diagrams
 Electrons that are in shells that are not
occupied by the nearest noble gas
element are called valence electrons.
 For Na, the 3s electrons are valence
electrons
 Valence electrons:
 Used to form chemical bonds
Orbital Diagrams
Example: Draw the orbital diagram for
potassium. How many valence electrons?
Know: Z = 19 so there are 19 electrons
1s
2s
2p
3s
3p
4s
Orbital Diagrams
Example: Draw the orbital diagram for Ti
(Z=22). How many valence electrons?
1s
2s
2p
3s
3d
3p
4s
Electron Configuration
 Drawing orbital diagrams gives information
not only about the orbitals that are/have
been filled but also about the number of
unpaired electrons.
 Orbital diagrams can be cumbersome!!
Electron Configuration
 A short-hand notation is commonly used in
place of orbital diagrams to describe the
electron configuration of an atom.
 Electron configuration:
a particular arrangement of electrons in
the orbitals of an atom
Electron Configuration
 The electron configuration tells the
number of electrons found in each
subshell.
 If there are three electrons in a 2p
subshell, we would write:
2p3
where the superscript (3) indicates the
number of electrons in that subshell
Electron Configuration
 The orbital diagram for an O atom:
1s
2s
2p
3s
The electron configuration for an O atom:
1s22s22p4
Electron Configuration
 To determine the electron configuration of
an atom (or ion) without first writing the
orbital diagram:
determine the number of electrons
present
add electrons to each subshell in the
correct filling order until all electrons
have been added
use the “diagonal” diagram to help
determine the filling order
Electron Configuration
Example: Write the electron configuration
of a Mn atom (Z = 25). How many valence e-?
1s22s22p63s23p64s23d5
Electron Configuration
Example: Write the electron configuration
of an O2- ion (Z = 8). With what noble gas is
this ion “isoelectric”?
An O2- ion has 8 protons and 10 electrons
1s22s22p6
Electron Configuration
For Sodium (Na)
11 electrons
1 more electron than the noble gas neon
1s2 2s2 2p6 3s1
Neon core
 ABBREVIATED or CORE notation
[Ne] 3s1
Electron Configuration
Write the full electron configuration for
 F
Then write the abbreviated or core notation
F
Electron Configuration
Write the full electron configuration for
 Rb
Then write the abbreviated or core notation
 Rb
Electron Configuration
Write the full electron configuration for
 Au
Then write the abbreviated or core notation
 Au
Electron Configuration
Write the abbreviated or core notation
 Es
Electron Configuration
To write the electron configuration using
core notation:
find the noble gas that comes before
the atom
determine how many additional
electrons must be added beyond
what the noble gas has
 Atomic number of atom minus
atomic number of noble gas
Electron Configuration
 To write the electron configuration using
core notation (cont):
determine the period number of the
element
this determines the value of n of
the s subshell to start with when
adding extra electrons
add electrons starting in the “n” s
subshell
Electron Configuration
Example: Write the core electron
configuration of Sr (Z = 38).
Previous noble gas: Kr (Z = 36)
Extra electrons: 38 - 36 = 2
Period number: 5
[Kr] 5s2
Electron Configuration
Example: Write the core electron
configuration of Br (Z = 35).
Previous noble gas: Ar (Z = 18)
Extra electrons: 35 - 18 = 17
Period number: 4
[Ar] 4s23d104p5
Orbital Diagrams
 Another useful periodic trend:
p block
d block
f block
Electron Configuration - Anomalies
Some irregularities
occur when there
are enough
electrons to halffill s and d orbitals
on a given row.
Electron Configuration - Anomalies
For instance, the
electron
configuration for
chromium is
[Ar] 4s1 3d5
rather than the
expected
[Ar] 4s2 3d4.
Isoelectronic Series
 When atoms ionize, they form ions with the
same number of electrons as the nearest (in
atomic number) noble gas.
Na = 1s22s22p63s1 = [Ne]3s1
Na+ = 1s22s22p6
= [Ne]
Cl = 1s22s22p63s23p5 = [Ne]3s23p5
Cl- = 1s22s22p63s23p6 = [Ar]
Isoelectronic Series
 N (7 e-):
 N3- (10 e-):
1s22s22p3
1s22s22p6 = [Ne]
 O (8 e-):
 O2- (10 e-):
1s22s22p4
1s22s22p6 = [Ne]
 F (9 e-):
1s22s22p5
 F- (10 e-):
1s22s22p6 = [Ne]
Isoelectronic Series
 Na (11 e-):
1s22s22p63s1
 Na+ (10 e-):
1s22s22p6 = [Ne]
 Mg (12 e-):
1s22s22p63s2
 Mg2+ (10 e-):
1s22s22p6 = [Ne]
 Al (13 e-):
 Al3+ (10 e-):
1s22s22p63s23p1
1s22s22p6 = [Ne]
1A
H
Ions of the highlighted
elements are
isoelectronic with Ne.
2A
Li Be
Na Mg
K
Rb
8A
3A
B
3B 4B 5B 6B 7B 8B
Ca Sc Ti
Sr Y
Zr
V
Cr Mn Fe
8B 8B 1B 2B
Co Ni
4A 5A 6A 7A He
C
N
O
F
Ne
Al Si
P
S
Cl
Ar
Se
Br Kr
Cu Zn Ga Ge As
Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb
Te I
Cs Ba La Hf
Ta W
Po At Rn
Fr Ra Ac Rf
Db Sg Bh Hs Mt
Re Os Ir
Pt
Au Hg Tl Pb Bi
Xe
Ce Pr
Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa
U
Np Pu Am Cm Bk Cf
Es Fm Md No Lr
Isoelectronic Series
 Isoelectronic: having the same number of
electrons
 N3-, O2-, F-, Ne, Na+, Mg2+, and Al3+ form an
isoelectronic series.
A group of atoms or ions that all contain
the same number of electrons
Isoelectronic Series
 Examples of isoelectronic series:
N3-, O2-, F-, Ne, Na+, Mg2+, Al3+
Se2-, Br-, Kr, Rb+, Sr2+, Y3+
Cr, Fe2+, and Co3+
Periodic Properties of Elements
 Chemical and physical properties of the
elements vary with their position in the
periodic table.
Atomic size
Ionization energy
Electronegativity
Metallic character
Periodic Properties--Atomic Size
 The relative size (radius) of an atom of
an element can be predicted by its
position in the periodic table.
 Trends
Within a group (column), the atomic
radius tends to increase from top to
bottom
Within a period (row), the atomic
radius tends to decrease as we move
from left to right
Periodic Table
Increasing size
Periodic Properties--Atomic Size
Increasing size
Lower “lefter” larger
Periodic Properties--Atomic Size
 Example: Which element would have the
larger atomic radius, Ar or Br?
Br should have the larger radius
more towards the bottom
more towards the left side
Periodic Properties
Ionization Energy
 The ease with which an electron can be
removed from an atom to form an ion is an
important indicator of its chemical behavior.
 Ionization energy: the minimum energy
required to remove an electron from the
ground state of an isolated gaseous atom or
ion.
Formation of cation (+) or more positively
charged cation
Periodic Properties
Ionization Energy
Ionization of Gaseous Sodium:
Na (g)  Na+ (g) + e As the ionization energy increases, it
becomes harder to remove an electron.
Periodic Properties
Ionization Energy
 Within each row, the ionization
energy increases from left to
right
 Its easiest to remove an
electron from an alkali metal
and hardest to remove one
from a noble gas.
 Within each group, the ionization
energy generally decreases from
top to bottom
 It’s easier to ionize K than Li.
Periodic Properties
Ionization Energy
Example: Which element has the higher
ionization energy, Br or Ca? Which one will
lose an electron easier?
Br has the higher ionization energy
further to the right
Ca will lose an electron easier because its
ionization energy is lower.
Periodic Properties
Electronegativity
 The ability of an atom to attract electrons in
a shared bond is called the electronegativity.
Cl (g) + e-  Cl- (g)
 The electron negativity increases as the
attraction between an atom and an electron
increases
more electronegativity = more likely to pull
an electron towards its nucleus
Periodic Properties
Electronegativity
 Trends:
Halogens have the highest electronegativity
Electronegativities increase moving from
the left toward the halogens.
Electronegativity decreases moving down a
column.
Noble gases have NO electronegativity
Periodic Properties
Electronegativity
 Trends:
Noble gases will not accept another electron.
To do so would require adding an electron
to a new electron shell (significantly higher
in energy)
Halogens have highest electronegativity.
Almost full valence shell
Metals have low electronegativity; metals
tend to lose electrons to form positive ions.
Closer to nucleus: higher electronegativity
Periodic Properties
Metallic Character
 Metals:
shiny luster
malleable and ductile
good conductors of heat and electricity
form cations
 Metallic character
increases from top to bottom
Increases from right to left