Transcript Energy Sublevels

Mr. Shields
Regents Chemistry
U06 L03
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Bohr Model
e- transitions from a higher energy levels to lower
energy levels release energy in the form of photons.
Bohr’s model correctly predicted hydrogen’s 4 visible
Lines in the emission spectra but incorrectly
predicted the emission spectra for all other atoms.
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Bohr Model
Obviously the Bohr model didn’t
Accurately describe the atom.
In the 1920’s a new model of the
Atom began to emerge.
It was known as the Quantum
Mechanical model.
Let’s look at some of the key concepts
That lead to this model
Niels Bohr
1885-1962
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The electron as a Wave
1924 – As a graduate student
de Broglie began considering
Bohr’s quantized atom.
If light could behave as a particle
Then why, de Broglie wondered,
couldn’t particles (like electrons)
behave as waves?
Louis Victor de Broglie
1892-1987
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The wave nature of an electron
De Broglie developed his idea and established a
Mathematical relationship that related wavelength
To mass & velocity
of a moving object.
h = 6.626 x 10
j sec
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In other words,
Anything that moves
behaves as a wave!
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Heisenberg
In 1913 Bohr showed how electrons were
quantized
In 1924 De Broglie showed how an electron
could behave as a wave.
In 1925 Werner Heisenberg established
what was known as matrix mechanics to
explain atomic behavior
Werner Heisenberg
1901- 1976
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Schrodinger
In 1926 Erwin Schrodinger proposed a
Wave Mechanics model of the atom.
The matrix and wave quantum model of
the atom divided scientists into two
camps.
Erwin Schrodinger
1887 - 1961
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The debate
In 1926 Werner Heisenberg began his job as Assistant to
Niels Bohr in Copenhagen. Later that year Schrodinger came
To debate the two alternative theories with Bohr.
Neither model was satisfactory but Schrodinger showed the
equivalence of the matrix and wave versions of Quantum
Mechanics.
After Bohr presented a statistical interpretation
of the wave function, these theories formed the
basis of what is now regarded as
quantum mechanics.
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Uncertainty Principle
1927 - Heisenberg proposed what he is most well
Known for … The Uncertainty Principle
“It is impossible to simultaneously know both the
Position and Velocity (a vector) of a particle at the
Same time.”
A photon has about the same energy as an electron.
A photon striking an electron causes the electron
To change both position and velocity. So an electrons
Position can not be specified with precision .
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Uncertainty principle
The consequence is that we can only give a
Probability of finding an electron in a given location.
Is the electron
*
Some location probabilities
are higher than others.
here
here
Or here?
*
*
Ans: yes to all but each has a
different probability
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The Quantum Mechanical model


Electrons behave as a wave or a particle with specific
allowable energy values – they are quantized.
Electrons are located in “probability regions” of space


These are known as atomic orbitals. They are not the
same as Bohr’s circular orbits!
An Orbital is defined as a region in space in which
there is a 90% probability of finding the electron
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Hydrogen’s 1st orbital;
electron probability plot
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The Quantum Mechanical model

Just like the Bohr Model there are Principal
quantum numbers (n)

n indicates distance from the nucleus


n Specifes major energy levels called the Principal
Energy Levels


As n increase distance from the nucleus increases
There are currently 7 principal energy levels
Principal energy levels contain Energy Sublevels
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Energy Sublevels
Each Principal Energy level contains 1 more sublevels than the
last principal energy level, beginning with 1 sublevel in the
1st Principal Energy Level
Sublevels are labeled s, p, d, and f (there are no others)
So…
n
1
2
3
4
5
6
7
number of
sublevels
1
2
3
4
4
4
4
Sublevel
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
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Energy Sublevels
Furthermore, within each energy sublevel there is anywhere
from 1 to 7 orbitals
Sublevel
# of Orbitals
s
p
d
f
1 s orbital
3 p orbitals
5 d orbitals
7 f orbitals
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Energy Sublevels
Electrons occupy the energy sublevels.
These energy sublevels refer to the 3
dimensional regions in space called atomic
Orbitals.
Remember: orbitals are defined as the region
In space in which there is a 90% probability
of finding the electron.
So what do these orbitals look like?
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The “s” orbital
90% probability
Inside circle
All “s” orbitals are spherical
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The s orbitals of the 1st three Energy Levels
n=1
n=2
2s
n=3
Note the growth in size as n increases
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The 3 “p” orbitals of the 2nd Principal Energy Level
The “Dumbbell” shaped p
Orbitals
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Calculating # of Orbitals
The number of orbitals in a given principle
energy level (n) is equal to n2
Remember “2n2 “ for calculating the # of e- in
each Bohr orbit ?
Well 2 is the max # of e- allowed in each orbital
hence 2n2 = max # of e- per principle quantum
number
So what do these 1s 2s and 2p atomic
orbitals look like as they surround the
nucleus ?
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Neon with fully occupied 1s, 2s, 2p orbitals
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All the orbitals of the
s, p, d, and f
sublevels
Only need to know these
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Orbital Energy
As the principal quantum number (n) increases, in
Other words as we move further from the nucleus,
The energy of the electrons in those principal
Energy levels increase;
Energy: 1 < 2 < 3 < 4 < 5 < 6 < 7
Also, within any principal energy level the energy of
The sublevels increase from s to f;
Energy: s < p < d < f
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Electron Orbital filling
This increasing energy sequence defines
into which Orbitals Electrons go as they
are added to the atom .
For example lets look at Rubidium …
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Sublevel
Energy increases
From s to f.
Some sublevels
With a lower n may
Actually be at a
Higher energy
Than some ein a higher n !
4f > 5s,5p,6s
4d > 5s
3d > 4s
We’ll look at this
“SWITCHING”
Again later
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