Transcript Site Name - chemdept

Orbital Diagrams
Element
Total
Electrons
H
1
He
2
Li
3
Be
4
1s
2s
Orbital Diagrams
Element
Total
Electrons
B
5
C
6
N
7
O
8
F
9
1s
2s
2p
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
 All other noble gases
 an octet of electrons in the outer shell
 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
Orbital Diagrams
For Sodium (Na)
 11 electrons
 1 more electron than the noble gas
neon
1s
2s
2p
Neon core
3s
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
 The ones lost to form cations
Orbital Diagrams
Example: Draw the orbital diagram for
potassium.
Know: Z = atomic number =
Orbital Diagrams
Example: Draw the orbital diagram for Ti.
Orbital Diagrams
 A useful periodic trend:
 For atoms in the 1st period, the
electrons are being added to the n=1
shell.
 For atoms in the 2nd period, the last
electrons are being added to the n=2
shell.
 Etc.
Orbital Diagrams
 Another useful periodic trend:
p block
d block
f block
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 order starting with the lowest
energy subshell until all electrons have
been added
 use the “diagonal” diagram to help
determine relative energy (i.e. filling
order)
Electron Configuration
Example: Write the electron configuration of
a Mn atom (Z = 25).
Electron Configuration
Example: Write the electron configuration of
an O2- ion (Z = 8).
Electron Configuration
Example: 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
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.
Electron Configuration
Example: Write the core electron configuration
of Br.
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+ =
 Cl
1s22s22p6
= [Ne]
= 1s22s22p63s23p5 = [Ne]3s23p5
 Cl- =
1s22s22p63s23p6 = [Ar]
Isoelectronic Series
 N (7 e-):
 N3- (10 e-):
 O (8 e-):
 O2- (10 e-):
 F (9 e-):
 F- (10 e-):
1s22s22p3
1s22s22p6 = [Ne]
1s22s22p4
1s22s22p6 = [Ne]
1s22s22p5
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-):
1s22s22p63s23p1
 Al3+ (10 e-):
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:
electrons
having the same number of
 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
 Size of Atom vs. Ion
 Size of Ions in Isoelectronic series
 Ionization energy
 Electron affinity
 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?
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.
Periodic Properties – Ion Size
 Trends to know:
 For ions in the same group (same charge),
size increases from top to bottom.
 Same trend as for the size of parent
atoms
 I- is larger than F-
Periodic Properties – Ion Size
 Trends to know:
 For an isoelectronic series of ions, the size
decreases with increasing atomic number.
 Na+ is smaller than O2-
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?
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
Periodic Properties
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.
Periodic Properties
Electron Affinity
 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)
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