bohr’s model of the atom. bohr’s model why don’t the electrons fall into the nucleus? e-...
TRANSCRIPT
Bohr’s Model of the Atom
Bohr’s Model
Why don’t the electrons fall into the nucleus? e- move like planets around the
sun. They move 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
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short Wavelength
Visible Light
Electromagnetic Spectrum
The line spectrum
Electricity passed through a gaseous element emits light at a certain wavelength
The colors can be seen when passed through a prism
Every gas has a unique pattern (color)
Further away from the nucleus means more energy.
There is no “in between” energy
First
Second
Third
Fourth
Fifth
Incr
easi
ng e
nerg
y }
Line spectrum of various elements
Electrons orbiting closest to the nucleus are said to be in the lowest energy state called the ground state
Atoms can absorb an amount of energy This promotes an electron to a higher energy
level called the excited state This energy level is unstable and so the
electron will fall back to its ground state When it does this, the excess energy will be
emitted as light
When the e- falls from one energy level to another, an amount of energy is emitted as light
This light emitted at specific wavelengths, which corresponds to our atomic spectra
Each atom will have different electron “jumps” therefore emitting different amounts of energy as light
This creates different line spectra for various elements
Let’s watch this video
http://www.mhhe.com/physsci/chemistry/chang7/esp/folder_structure/pe/m1/s3/
More videos (view OYO)
http://www.mhhe.com/physsci/astronomy/applets/Bohr/applet_files/Bohr.html
http://highered.mcgraw-hill.com/sites/0072482621/student_view0/interactives.html#
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 e-.
His theory could only explain hydrogen and not the more complex atoms
The Quantum Mechanical Model
Energy is quantized – meaning 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.
An equation has been developed 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 equation describes several shapes.
These are called atomic orbitals Orbitals are regions where there is a
high probability of finding an e-
S sublevel
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 sublevel
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
The 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 has seven different shapes (orbitals)
2 electrons per orbital
Summary
Sublevel # of Orbitals
# e- in sublevel
Starts in what energy level
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- e- enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2 e-per orbital with different spins
Hund’s Rule- When e- occupy orbitals of equal energy they don’t pair up until they have to .
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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
Electron Configuration for phosphorus
The first two electrons go into the 1s orbital
Notice the opposite spins
only 13 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 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
Write these electron onfigurations
Titanium
Vanadium
Chromium
Copper
Electron Configurations
Titanium - 22 electrons 1s2
2s2
2p6
3s2
3p6
4s2
3d2
Vanadium - 23 electrons 1s2
2s2
2p6
3s2
3p6
4s2
3d3
Chromium - 24 electrons 1s2
2s2
2p6
3s2
3p6
4s2
3d4
Expected. But, this is wrong!!
1s2
2s2
2p6
3s2
3p6
4s1
3d5
Copper – 29 electrons 1s2
2s2
2p6
3s2
3p6
4s2
3d9
Actual configuration is
1s2
2s2
2p6
3s2
3p6
4s1
3d10
Why are chromium and copper configurations different? This gives two half filled orbitals for chromium and one
completely full and one half filled orbital for copper.
Slightly lower in energy. Remember these exceptions
Practice
Element 1s 2s 2px 2py 2pz 3s Electron ConfigurationH
HeLiCNOF
NeNa
Electron Configuration for Some Selected ElementsOrbital filling