Transcript Electron Configuration

The Elements
CH. 2
atomic models
electronic configuration
oxidation numbers
The Models of the Atom
Line Spectra
What is causing the light to appear when the
atom is excited by electricity?
Why does the line spectrum of hydrogen have
very distinct lines of color and not the
complete rainbow, which includes all
colors?
The Atomic Energy and the Bohr
Model of the Atom
h
The Quantum Theory
• The main idea that comes from the
Quantum Mechanical Theory and the
Schrodinger Equation is Quantum numbers
that define the surfaces of the electron
densities in which the electrons reside.
• There are 3 main quantum numbers, n, l and
ml, and a fourth quantum number, ms
Principle quantum number, n
Schrödinger’s equation provides a
number of possible functions that define
the shape of the electron cloud about the
nucleus, these shapes are the orbitals.
The functions can be arranged in sets and
subsets. Orbitals are classified in three
ways:
An electron shell is a collection of orbitals, identified by
an integer, n, which ranges in value infinitely up from
1. Electrons with the same value are in the same shell.
In chemistry, important shells have n values between 1
and 7.
Angular Momentum, l
Subshells are groups of orbitals within an electron shell.
They are identified by the letters s, p, d, f and through
the alphabet beginning with g. In chemical systems,
only s, p, d, and f subshells are important. The number
of subshells within an electron shell is equal to the n
value of the shell.
Magnetic, ml
Individual orbitals are identified by their direction in
space using Cartesian coordinates. The number of
orbitals within a subshell depends on the subshell type:
s subshells have a single orbital, p subshells have 3
orbitals, d subshells 5 and f subshells 7 orbitals.
Values of n, l and ml
Example
n value
l value
n = 1, 2,… l = 0, 1, .. n-1
n=1
l=0
n =2
l=0
l=1
n =3
l=0
l=1
l=2
ml value
ml = 0, 1, …l
ml = 0
ml = 0
ml = +1, 0, -1
ml = 0
ml = +1, 0, -1
ml = +2,+1,0,-1,-2
The orbital shape and l
n value l value
n=1
l=0
n =2
l=0
l=1
n =3
l=0
l=1
l=2
n =4
l=0
l=1
l=2
l=3
Orbital designation
1s
2s
2px, 2py, 2pz
3s
3p
2
2 2
3dxy, 3dz , 3dyz, 3dxz, 3dx -y
4s
4p
4d
4f
The Energy Levels of the orbitals
In hydrogen the energy of
the orbitals and the order
in which they fill are
dependant only upon
their principle quantum
number; however, in
atoms with more
electrons the order
changes.
Energy levels in atoms of more
than one electron.
Notice the shifting of the
energy levels, as n
increases, energy
increases, as l increases,
energy increases, but
mixing of the energy
levels still occur.
Orbitals
10s
9s
9p
9d
9f
9g
9h
9i
9j
8s
8p
8d
8f
8g
8h
8i
8j
7s
7p
7d
7f
7g
7h
7i
6s
6p
6d
6f
6g
6h
5s
5p
5d
5f
5g
4s
4p
4d
4f
3s
3p
3d
2s
2p
1s
9k
Increasing
Energy
• Electrons fill different orbitals in a subshell until the
subshell is half - filled. (one in each orbital before
pairing begins)
• If an atom has a filled valence shell, the atom is
more stable.
• Each electron or the space it occupies can be defined
by its four quantum numbers (n, l, ml, ms)
• No two electrons can occupy exactly the same space
or have the same four quantum numbers.
• Some electron configurations defy expectations
created by the periodic table, particularly for heavier
elements.
Periodic Blocks
Why do we have those
rows at the bottom?
H
He
Li Be
B
Na Mg
Al Si P
K Ca Sc
Ti
Rb Sr Y
C
N
O
F Ne
S Cl Ar
V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I
Xe
Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
This arrangement takes too
much space and is hard to read.
Hund’s Rule
A full energy level is very
stable, of lowest energy, a
half filled shell is almost
as stable as a fully filled
shell and a partially filled
shell is the least stable of
highest energy. The
orbitals will go from
partially filled to half
filled by putting one
electron in each orbital
before pairing the spins.
Electron Configurations
• Complete electron configurations
• noble gas electron configurations showing
the noble gas and the remaining valence
electrons
• valence electron configuration
Fe
4s
3d
+2
Fe
4s
3d
+3
Fe
4s
3d
Rules for assigning oxidation numbers
The sum of the oxidation number of all atoms
must equal the net charge of the species.
In compounds:
Group IA are +1.
Group IIA are +2.
B and Al are +3, and F is -1.
Hydrogen is +1 except when combined with a metal. Then
it is -1.
Oxygen is -2 except for peroxides and superoxides.
Elements in their elemental state have an oxidation
number of zero.
Li
Be
B
+1
+2
+3
H
He
C
N
+4 +5 +4
-2 +3 +2
-4 +1 -3
O
F
-1
-2
-1
S
Cl
Ne
+1
Na Mg
Al
+1
+3
+2
K
Ca Sc
+1
+2
Ti
V
3+
Rb Sr
Y
+1
+3
Mo Tc Ru Rh Pd
Nb
Zr
+6
+7 +8 +6 +4
+4 Ag Cd
+2
+4
Cs Ba Lu
Hf
Ta
+1
+4
+5
+2
+3
Fr Ra Lr
+1
+2
+3
+4
-4
P
+5 +6 +4 +7 +5
+3
+2 +3 +1
-3
-2
-1
Cr Mn Fe Co
As Se Br
Cu
Ge
Ni
Zn Ga
+6 +7 +6
5+
6+
+5
+4 +5 +4
+3
+3
+3
+3 +4 +3
+2
+2
+2
+2
+2
+5
+4
Si
+4
+3
+6
+4
+3
+2
+2
+2
+1
In
+4
-4
3+
3-
+3
+4
+2
+5
+3
-3
+2
W Re Os
Ir
Pt Au Hg
Tl
Pb
Bi
+6
+4
+4
+3
+4
+2
+3
+1
+4
+2
+5
+3
+3
+1
+2
+1
4+
2-
+1
-1
Sn Sb Te
+3
+2
+8
+6
+2
+3
+4
+3
+7
+6
+4
+1
+2
I
Ar
Kr
+4
+2
Xe
+6 +7 +5 +6
+4
+1
+4
-2
-1
+2
Po
+2
At Rn
Common oxidation numbers
-1
Oxidation numbers and the periodic table
Some observed trends in compounds.
Metals have positive oxidation numbers.
Transition metals typically have more than
one oxidation number.
Nonmetals and semimetals have both positive
and negative oxidation numbers.
No element exists in a compound with an
oxidation number greater than +8.
The most negative oxidation numbers equals 8
- the group number
Electron Configuration Activities
• Simple Atomic Structure
• Atomic Structure - Electron Configuration
• Structured Curriculum Lesson Plan