# Orbital Diagrams and Electron Configuration by davebusters

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```									            Orbital Diagrams and
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!!

1
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

2
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

3
Electron Configuration
• The orbital diagram for an O atom:

1s    2s      2p         3s

The electron configuration for an O atom:

1s22s22p4

4
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
• use the “diagonal” diagram to help determine
the filling order

5
Electron Configuration
Example: Write the electron configuration
of a Mn atom (Z = 25).

1s2 2s22p6 3s23p6 4s2 3d5

6
Electron Configuration
Example: Write the electron configuration
of an O2- ion (Z = 8).
An O2- ion has 8 protons and 10 electrons

1s22s22p6

7
Electron Configuration
Write the electron configuration of a krypton
atom (Z = 36).
1s22s22p63s23p64s23d104p6

This is the Kr “core”  [Kr]

• The noble gas “core” can be used to write the
electron configuration of an element using
core notation:
noble gas “core” + valence electrons

8
Core notation
To write the electron configuration using the core notation:
•   Find the Noble Gas that comes before the atom.
•   Determine how many additional electrons must be
added beyond what that noble gas has.
(= Atomic number of atom minus atomic number of
noble gas)
•   Determine the period that element is in. (This
determines the value of n of the s subshell to start with
•   Add electrons starting in that “n” subshell

9
Electron Configuration
Write the core electron configuration of Sr
(Z = 38).
Previous noble gas: Kr (Z = 36)
Extra electrons: 38 (e of Sr) - 36 = 2
Period number of Sr: 5
So: Kr core plus 2 extra e- starting in 5s
[Kr] 5s2

10
Electron Configuration
Write the core electron configuration of Br
(Z = 35).
Previous noble gas: Ar (Z = 18)
Extra electrons: 35 - 18 = 17
Period number: 4
So: Ar core plus 17 extra e- starting with 4s
[Ar] 4s23d104p5

11
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]

12
Isoelectronic Series

• N (7 e-):      1s22s22p3
 N3- (10 e-):   1s22s22p6 = [Ne]

• O (8 e-):      1s22s22p4
 O2- (10 e-):   1s22s22p6 = [Ne]

• F (9 e-):      1s22s22p5
 F- (10 e-):    1s22s22p6 = [Ne]
13
Isoelectronic Series

• Na (11 e-):     1s22s22p63s1
 Na+ (10 e-):    1s22s22p6 = [Ne]

• Mg (12 e-):     1s22s22p63s2
 Mg2+ (10 e-):   1s22s22p6 = [Ne]

• Al (13 e-):     1s22s22p63s23p1
 Al3+ (10 e-):   1s22s22p6 = [Ne]

14
1A
Ions of the highlighted                                   8A
H    2A
elements are
3A   4A 5A 6A 7A He
isoelectronic with Ne.
Li Be                                           B   C   N   O    F    Ne

Na Mg                                           Al Si   P   S    Cl   Ar
3B 4B 5B 6B 7B 8B      8B 8B 1B 2B

K    Ca Sc Ti     V   Cr Mn Fe   Co Ni    Cu Zn Ga Ge As    Se   Br Kr

Rb   Sr Y    Zr   Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb          Te I      Xe

Cs Ba La Hf       Ta W    Re Os Ir   Pt   Au Hg Tl Pb Bi    Po At Rn

Fr Ra Ac Rf       Db Sg Bh Hs Mt

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
15
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

16
Isoelectronic Series
• Examples of isoelectronic series:
– N3-, O2-, F-, Ne, Na+, Mg2+, Al3+

– Se2-, Br-, Kr, Rb+, Sr2+, Y3+

– Cr, Fe2+, and Co3+

17
Periodic Properties of Elements
• Chemical and physical properties of the
elements vary with their position in the periodic
table.
– Atomic size
– Size of Atom vs. Ion
– Size of Ions in Isoelectronic series
– Ionization energy
– Electron affinity
– Metallic character
18
What Is the Size of an Atom?
The bonding atomic
one-half of the
distance between
covalently bonded
nuclei.

19
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

20
Sizes of Atoms
Bonding atomic
…decrease from left to
right across a row

…increase from top to
bottom of a column

21
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

22
Periodic Properties – Atom vs. Ion Size

• Trends to know:
– Cations (+) are smaller than their
parent atoms.
• Electrons are removed from the outer
shell.
– Anions (-) are larger than their
parent atoms.
• Electron-electron repulsion causes the
electrons to spread out more in space.
23
Sizes of Ions - Trends
• Ions increase in size
as you go down a
column.

24
Sizes of Ions - Trends

• In an isoelectronic series, ions have the same
number of electrons.
• Ionic size decreases with an increasing nuclear
charge.

25
Ionization Energy

• The ionization energy is the amount of
energy required to remove an electron
from the ground state of a gaseous
atom or ion to form a cation or more
positively charged cation.
– The first ionization energy is that energy
required to remove first electron.
– The second ionization energy is that
energy required to remove second
electron, etc.
26
Ionization of Gaseous Sodium:

Na (g)  Na+ (g) + e-

• As the ionization energy increases, it
becomes harder to remove an electron.

27
Ionization Energy
• It requires more energy to remove each
successive electron.
• When all valence electrons have been removed,
the ionization energy takes a quantum leap.

28
Trends in First Ionization
Energies
• As one goes down a
column, less energy
is required to remove
the first electron.
– valence electrons are
farther from the
nucleus.
Within each row, the
ionization energy
increases from left to
right
29
Ionization Energy
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.

30
Periodic Properties
Electron Affinity
• The energy change that occurs when an electron is
added to a gaseous atom is called the electron affinity.

Cl (g) + e-  Cl- (g)

• The electron affinity becomes increasingly negative as
the attraction between an atom and an electron
increases
– more negative electron affinity = more likely to gain an
electron and form an anion

31
Electron Affinity
• Trends:
– Halogens have the most negative electron affinities.

– Electron affinities become increasing negative moving
from the left toward the halogens.

– Electron affinities do not change significantly within a
group.
– Noble gases will not accept another electron.
• To do so would require adding an electron to a
new electron shell (significantly higher in energy)

32
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

33
Properties of Metal, Nonmetals,
and Metalloids

34
Metals versus Nonmetals

Differences between metals and nonmetals tend
to revolve around these properties.

35
Metals versus Nonmetals
• Metals tend to form cations.
• Nonmetals tend to form anions.

36
Metals
They tend to be
lustrous, malleable,
ductile, and good
conductors of heat
and electricity.

37
Metals
• Compounds formed
between metals and
nonmetals tend to be
ionic.
• Metal oxides tend to
be basic.

38
Nonmetals
• These are dull, brittle
substances that are
poor conductors of
heat and electricity.
• They tend to gain
electrons in reactions
with metals to acquire
a noble gas
configuration.

39
Nonmetals
• Substances
containing only
nonmetals are
molecular
compounds.
• Most nonmetal oxides
are acidic.

40
Metalloids
• These have some
characteristics of
metals and some of
nonmetals.
• For instance, silicon
looks shiny, but is
brittle and fairly poor
conductor.

41

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