Bohr model and electron configuration Prof. Mohammed Gulam Ahamad CE2100-Semiconductors.
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Transcript of Bohr model and electron configuration Prof. Mohammed Gulam Ahamad CE2100-Semiconductors.
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 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)
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
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
The line spectrum
electricity passed through a gaseous element emits light at a certain wavelengthCan be seen when passed through a prismEvery gas has a unique pattern (color)
Bohr’s Triumph
His theory helped to explain periodic law
Halogens are so reactive because it has one e- less 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
Further away from the nucleus means more energy.
There is no “in between” energy
Energy LevelsFirst
Second
Third
Fourth
Fifth
Incr
easi
ng e
nerg
y }
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 D sublevel contains 5 D orbitalsThe D sublevel starts in the 3rd 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
s
p
d
f
# of shapes (orbitals)
Max # of electrons
1 2 1
3 6 2
5 10 3
7 14 4
Sublevel
Starts at energy level
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 Levelonly s sublevel (1 s orbital)only 2 electrons1s2
Second Energy Levels and p sublevels (s and p orbitals are available)2 in s, 6 in p2s22p6
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
Electron ConfigurationHund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to .
The first two electrons go into the 1s orbital
Notice the opposite spins
only 13 moreIncr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
The next electrons go into the 2s orbital
only 11 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 2p orbital
• only 5 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 3s orbital
• only 3 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• 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
1s22s22p63s23p64s23d2
Vanadium - 23 electrons 1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 is expected
But this is wrong!!
Chromium is actually1s22s22p63s23p64s13d5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.
Copper’s electron configurationCopper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.Remember these exceptions