Download ORBITALS - BROCHEM

BOHR’S MODEL
Why don’t the electrons fall into the
nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level
from another.
BOHR’S MODEL
Nucleus
Electron
Orbit
Energy Levels
BOHR POSTULATED THAT:
Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away from
the nucleus
An atom with maximum number of electrons in the
outermost orbital energy level is stable (unreactive)
ELECTROMAGNETIC
RADIATION
ELECTROMAGNETIC RADIATION.
ELECTROMAGNETIC
RADIATION
Most subatomic particles behave as PARTICLES and obey the physics of
waves.
ELECTROMAGNETIC
RADIATION
wavelength
Visible light
Amplitude
wavelength
Ultaviolet radiation
Node
ELECTROMAGNETIC
RADIATION
Waves have a frequency
Use the Greek letter “nu”,
are “cycles per sec”
, for frequency, and units
  = c
All radiation:
•
where c = velocity of light = 3.00 x 108 m/sec
HOW DID HE DEVELOP HIS THEORY?
•He used mathematics to explain the visible
spectrum of hydrogen gas
•http://www.mhhe.com/physsci/chemistry/
essentialchemistry/flash/linesp16.swf
High
Low
energy
energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Wavelength
Visible Light
THE LINE SPECTRUM
•electricity passed
through a gaseous
element emits light at
a certain wavelength
•Can be seen when
passed through a
prism
•Every gas has a
unique pattern (color)
LINE SPECTRUM OF VARIOUS ELEMENTS
BOHR’S TRIUMPH
His theory helped to explain periodic law
Halogens are so reactive because it has one eless than a full outer orbital
Alkali metals are also reactive because they
have only one e- in outer orbital
DRAWBACK
•Bohr’s theory did not
explain or show the
shape or the path
traveled by the
electrons.
•His theory could only
explain hydrogen and
not the more complex
atoms
ATOMIC LINE EMISSION
SPECTRA AND NIELS
BOHR
Niels Bohr
(1885-1962)
Bohr’s greatest contribution to
science was in building a simple
model of the atom. It was based
on an understanding of the LINE
EMISSION SPECTRA of excited
atoms.
Problem is that the model only
works for H
ATOMIC SPECTRA
One view of atomic structure in early 20th century was that an
electron (e-) traveled about the nucleus in an orbit.
ATOMIC SPECTRA AND
BOHR
Bohr said classical view is wrong.
Need a new theory — now called
QUANTUM or WAVE MECHANICS.
e- can only exist in certain discrete orbits
e- is restricted to QUANTIZED energy state
(quanta = bundles of energy)
QUANTUM OR WAVE
MECHANICS
Schrodinger applied idea of ebehaving as a wave to the
problem of electrons in atoms.
He developed the WAVE
EQUATION
Solution gives set of math
expressions called WAVE
E. Schrodinger
FUNCTIONS, 
1887-1961
Each describes an allowed energy
state of an e-
HEISENBERG UNCERTAINTY
PRINCIPLE
W. Heisenberg
1901-1976
Problem of defining nature
of electrons in atoms
solved by W. Heisenberg.
Cannot simultaneously
define the position and
momentum (= m•v) of an
electron.
We define e- energy exactly
but accept limitation that
we do not know exact
position.
ARRANGEMENT OF
ELECTRONS
IN
ATOMS
Electrons in atoms are arranged as
LEVELS (n)
SUBLEVELS (l)
ORBITALS (ml)
QUANTUM NUMBERS
The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe
the location of an electron within an atom or ion
n (principal)
---> energy level
l (orbital) ---> shape of orbital
ml (magnetic) ---> designates a particular
suborbital
The fourth quantum number is not derived from the wave function
s (spin)
---> spin of the electron
(clockwise or counterclockwise: ½ or – ½)
QUANTUM NUMBERS
So… if two electrons are in the same place at the
same time, they must be repelling, so at least the
spin quantum number is different!
The Pauli Exclusion Principle says that no two
electrons within an atom (or ion) can have the same
four quantum numbers.
If two electrons are in the same energy level, the
same sublevel, and the same orbital, they must
repel.
Think of the 4 quantum numbers as the address of an
electron… Country > State > City > Street
ENERGY LEVELS
Increasing energy
Fifth
Fourth
Further away from the
nucleus means more
energy.
There is no “in between”
energy
Energy Levels
Third
Second
First
THE QUANTUM MECHANICAL MODEL
Energy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move
from one energy level to another.
Since the energy of an atom is never “in between”
there must be a quantum leap in energy.
Schrödinger derived an equation that described the
energy and position of the electrons in an atom
ATOMIC ORBITALS
Principal Quantum Number (n) = the energy level
of the electron.
Within each energy level the complex math of
Schrödinger's equation describes several
shapes.
These are called atomic orbitals
Regions where there is a high probability of
finding an electron
S ORBITALS
1 s orbital for
every energy level
1s
2s 3s
Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals
P ORBITALS
Start at the second energy level
3 different directions
3 different shapes
Each orbital can hold 2 electrons
The p Sublevel has 3 p
orbitals
THE D SUBLEVEL CONTAINS 5 D
ORBITALS
rd
•The D sublevel starts in the 3 energy level
•5 different shapes (orbitals)
•Each orbital can hold 2 electrons
THE F SUBLEVEL HAS 7 F ORBITALS
•The F sublevel starts in the fourth energy level
•The F sublevel has seven different shapes (orbitals)
•2 electrons per orbital
SUMMARY
Starts at
energy
level
Sublevel
# of shapes
(orbitals)
Max # of
electrons
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
ELECTRON CONFIGURATIONS
The way electrons are arranged in atoms.
Aufbau principle- electrons enter the
lowest energy first.
This causes difficulties because of the
overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2
electrons per orbital - different spins
ELECTRON CONFIGURATIONS
First Energy Level
only s sublevel (1 s orbital)
only 2 electrons
1s2
Second Energy Level
s and p sublevels (s and p orbitals are available)
2 in s, 6 in p
2s22p6
8 total electrons
Third energy level
s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s24p64d104f14
32 total electrons
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
ELECTRON CONFIGURATION
Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until
they have to .
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p The first to electrons go
into the 1s orbital
2p
Notice the opposite spins
only 13 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p The next electrons go into
the 2s orbital
2p
only 11 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
ORBITALS FILL IN ORDER
Lowest energy to higher energy.
Adding electrons can change the energy of the orbital.
Half filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order
WRITE THESE ELECTRON CONFIGURATIONS
Titanium - 22 electrons
2 2 6 2 6 2 2
1s 2s 2p 3s 3p 4s 3d
2 2 6 2 6 2 3
Vanadium - 23 electrons 1s 2s 2p 3s 3p 4s 3d
Chromium - 24 electrons
2 2 6 2 6 2 4 is expected
1s 2s 2p 3s 3p 4s 3d
But this is wrong!!
CHROMIUM IS ACTUALLY
2 2 6 2 6 1 5
1s 2s 2p 3s 3p 4s 3d
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.
COPPER’S ELECTRON CONFIGURATION
Copper has 29 electrons so we expect
2 2 6 2 6 2 9
1s 2s 2p 3s 3p 4s 3d
But the actual configuration is
2 2 6 2 6 1 10
1s 2s 2p 3s 3p 4s 3d
This gives one filled orbital and one half filled orbital.
Remember these exceptions
PRACTICE
Time to practice on your own filling up electron configurations.
Do electron configurations for the first 20 elements on the periodic table.