CHAPTER 5 - ELECTRONS IN ATOMS
Note Taking Guide: Episode 303
Bohr's Energy Levels
Electrons in CERTAIN ENERGY LEVELS
• LOW energy levels: CLOSER to NUCLEUS
• HIGH energy levels: FARTHER from NUCLEUS
• Ground State: ELECTRON in LOWEST ENERGY LEVEL
possible
Excited Atom
• Atom has ABSORBED ENERGY.
• EXCITED state is UNSTABLE.
• ATOM soon EMITS same amount of ENERGY ABSORBED.
• ENERGY seen as VISIBLE LIGHT.
Wave Description of Light
Wavelength (): DISTANCE between CORRESPONDING POINTS
on ADJACENT waves
Frequency ( ): the NUMBER of WAVES passing a given
POINT in a given TIME
c =  or f
c = 3.0 X 108 m/s: speed of LIGHT
Sample problem #1: What is the frequency of light if the wavelength
is 6.0 x 10 -7 m?
 = 6.0 x 10 -7 m
c = 3.0 X 108 m/s
c = 
= c/ 3.0 X 108 m/s = 5.0 X 1014/s = 5.0 X 1014 Hz
6.0 x 10 -7 m
Sample problem #2: What is the wavelength of light if its frequency
is 5.0 x 1014 Hz?
 = 5.0 x 1014 Hz
c = 3.0 X 108 m/s
c = 
 = c/ 3.0 X 108 m/s = 6.0 x 10 -7 m
5.0 x 1014 Hz(1/s)
Particle Description of Light
ENERGY exists as PARTICLES called QUANTA
E = h
The Modern View of Light
LIGHT has a DUAL NATURE.
Light may EXIST as a WAVE.
Light may BEHAVE as a STREAM of PARTICLES called QUANTA
or PHOTONS.
Spectroscopy
SPECTRAL lines represent
ENERGY RELEASED as
ELECTRONS returns to
LOWER ENERGY LEVELS.
SPECTRAL lines IDENTIFY an ELEMENT.
Called the BRIGHT LINE SPECTRA of an ELEMENT.
Orbital -REGION of SPACE where an ELECTRON is LIKELY to
be FOUND
The Chemistry Quiz
CR1. A CR2. A
1.C
2.A
3.B
4.D
5.A
The Rutherford-Bohr atomic model is called the PLANETARY
model. RUTHERFORD said that the atom had a very SMALL,
DENSE nucleus. Niels Bohr pictured the electrons ORBITING
around the NUCLEUS. These electrons can jump from one
ENERGY LEVEL to another.
An ENERGY LEVEL of an electron is the region around the
nucleus where the electron is likely to be moving.
When e- fall back to lower energy states they can
give off their energy in the form of LIGHT energy.
http://www.visionlearning.com/library/module_viewer.php?mid=51&l=&c3=
http://www.dlt.ncssm.edu/core/c8.htm
Light is a form of ELECTROMAGNET energy.
Electromagnetic energy travels in WAVES.
Frequency is measured in HERTZ(Hz)or s-1= (1/s)
http://chemmovies.unl.edu/ChemAnime/DEFLITD/DEFLITD.html
Rabbits Mate In Very Unusual eXpensive Gardens
Electromagnetic energy is also called electromagnetic
RADIATION.
RADIO WAVES, MICROWAVES, INFRARED RAYS, VISIBLE
LIGHT, ULTRAVIOLET RAYS, X-RAYS, and GAMMA RAYS
are types of electromagnetic energy.
electromagnetic spectrum
http://csep10.phys.utk.edu/astr162/lect/light/spectrum.html
http://lectureonline.cl.msu.edu/~mmp/applist/Spectrum/s.htm
http://csep10.phys.utk.edu/astr162/lect/light/absorption.html
http://www.darvill.clara.net/emag/emagorder.htm
Long wavelength
Short
wavelength
Low frequency
High frequency
Low energy
High energy
The higher the frequency the SHORTER the wavelength.
Red light has a LOW frequency and LONG wavelength.
Blue light has HIGH frequency and SHORT wavelength.
Each kind of atom gives off its own SPECTRUM of
electromagnetic energy.
A SPECTRUM is the range of wavelengths emitted or absorbed
by excited atoms.
EMISSION spectra show the wavelengths given off when
excited ELECTRONS fall back to LOWER energy levels. A
QUANTUM of energy is the amount of energy required to move
an electron from one energy level to another.
ABSORPTION spectra shows the
wavelengths absorbed when ELECTRONS
are excited jump to HIGHER energy levels.
The element HELIUM was discovered in the
emission spectrum of the sun.
emission & absorption spectra
http://jersey.uoregon.edu/vlab/elements/Elements.html
PLANCK'S HYPOTHESIS
PLANCK developed the quantum theory.
The quantum theory states that energy is given off is packets
called QUANTA. The amount of energy is equal to
PLANCK'S CONSTANT(h) x FREQUENCY()
E = h
-34
h = PLANCK'S CONSTANT= 6.626 x 10 J/Hz
Ex: What is the energy of a photon from the violet portion of the
rainbow if it has a frequency of 7.23 x 1014 s-1.
E = h6.626 x 10-34 J | 7.23 x 1014 s-1| 1 Hz = 4.79 x 10-19 J
Hz
|
| 1 s-1
Each line of the SPECTRUM of an element corresponds to an
ENERGY change when electrons fall to lower energy levels.
In the PHOTOELECTRIC effect, electrons called PHOTO
electrons are ejected by metals when light shines on them.
EINSTEIN explained the photoelectric effect by saying that the
electrons absorbed photons of light energy. The energy of the
emitted electrons is proportional to the FREQUENCY of the light
that shines on them. A greater intensity of light makes MORE
electrons be emitted but doesn’t change their ENERGY.
photoelectric effect
http://www.wellesley.edu/Chemistry/Flick/peeffect5.html
http://www.lewport.wnyric.org/mgagnon/Photoelectric_Effect/photoelectriceffect1.html
http://www.usd.edu/phys/courses/phys431/notes/notes5g/photoelectric.html
deBROGLIE -said that particles have characteristics of waves.
 = h/mv
h= PLANCK'S CONSTANT m = MASS
v = VELOCITY


 = WAVELENGTH
As the momentum (mv) of an object increases its
WAVELENGTH decreases. Even large objects have a
WAVELENGTH but it’s too SMALL to be measured.
Wave property of objects
http://www.wwnorton.com/chemistry/overview/ch3.htm
The wave-particles duality of nature means that PARTICLES
can act as waves and WAVES can act as particles.
The HEISENBERG UNCERTAINTY principle states that the
exact position and momentum of an object can not be
determined at the same time. In order to see and object it must
be hit with a PHOTON of light which causes a change in the
VELOCITY of the object.
SCHRODINGER developed an equation to describe the wavelike behavior of the electron.
MAX BORN used Schrodinger's equation to find the probability
of finding an electron at any specific spot around the atom.
The location of an electron around the atom can best be
described as a CLOUD.
QUANTUM NUMBERS
QUANTUM NUMBERS represent different electron energy
states and are used to describe the electron.
QUANTUM mechanics is used to describe the behavior of
extremely small particles traveling at velocities near the speed
of light.
NEWTONIAN mechanics is used to describe the behavior of
larger objects at slower velocities
Quantum Numbers
- n,l,m,s
-Used to DESCRIBE an ELECTRON in an ATOM
n
-PRINCIPAL QUANTUM NUMBER
-Represents MAIN energy level of ELECTRON
- shows the SIZE of the electron cloud
-MAXIMUM # of ELECTRONS in an ENERGY LEVEL = 2n2
Example: What is the maximum number of electrons that can be in
the 7 main energy level? 2(7)2 = 98
l
-The SECOND QUANTUM NUMBER
- Describes the ORBITAL SHAPE within an ENERGY LEVEL
- NUMBER of orbital SHAPES possible in ENERGY LEVEL=n
- describes the SUBLEVEL.
Orbital Shapes
designated s, p, d, f
s
s, p
s, p, d
s, p, d, f
How many electrons can each sublevel hold?
s = 1 orbital x 2 e-/orbital = 2 ep = 3 orbitals x 2 e-/orbital = 6 ed = 5 orbitals x 2 e-/orbital = 10 ef = 7 orbitals x 2 e-/orbital = 14 e-
Energy
Level (n)
Principal
Quantum #
1
2
3
4
Sublevel
(type of
orbital) 2nd
quantum #
s
s
p
s
p
d
s
p
d
f
# of Orbitals
in Sublevel
3rd quantum
#
1
1
3
1
3
5
1
3
5
7
# of e- in
sublevel
2
2
6
2
6
10
2
6
10
14
Total # of
e- in
energy
level
2
8
18
32
m
THIRD QUANTUM NUMBER
ORIENTATION of ELECTRON in ORBITAL
s
FOURTH QUANTUM NUMBER (+1/2 or -1/2)
SPIN of ELECTRON in ORBITAL
Ground State: LOWEST energy arrangement of ELECTRONS
Diagonal Rule
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
7s2
7p6
7d10
7f14
8s2
8p6
8d10
Examples—
hydrogen 1s1
Sodium 1s2 2s2 2p6 3s1
Lithium 1s2 2s1
Nitrogen 1s2 2s2 2p3
Orbital Notation
Examples—
H (Z = 1)
hydrogen

1s
2s
2p
3s
3p
3s
3p
4s
3d
nitrogen
(Z = 7)


1s
2s


2p

4s
3d
Hund's Rule:
ORBITALS of EQUAL ENERGY are each FILLED by one
ELECTRON with PARALLEL spin before any ORBITAL is
occupied by a SECOND ELECTRON.
Pauli Exclusion Principle:
No two ELECTRONS in the SAME ATOM can have the SAME
SET of FOUR QUANTUM NUMBERS. If the 2 electrons are in
the same, energy level, sublevel and orbital they must have
opposite SPINS.
The Chemistry Quiz
CR1. B CR2. C 1. B 2. B 3. D 4. A 5. C
http://micro.magnet.fsu.edu/electromag/java/atomicorbitals/
DEGENERATE sublevels (p, d, f) have orbitals with the
same ENERGY but with different ORIENTATION in space.
http://www.wwnorton.com/chemistry/overview/ch3.htm
Electron configuration
http://www.wwnorton.com/chemistry/overview/ch3.htm
ELECTRON CONFIGURATION
The ATOMIC NUMBER (Z) represents the number of protons in
an atom. If an atom is NEUTRAL number of positive protons =
the number of negative ELECTRONS. The ATOMIC NUMBER
also gives the number of electrons in a neutral atom. The sum
of the superscripts in the electron configuration will be equal to
the ATOMIC NUMBER. Three rules are used to write electron
configurations.
1) AUFBAU principle states that the electrons enter orbitals of
LOWEST energy first.
2) The PAULI EXCLUSION principle.
3) HUND’S rule
Order of filling energy levels, sublevels and orbitals
7p
6d
5f
7s
___ ___ ___
___ ___ ___ ___ ___
___ ___ ___ ___ ___ ___ ___
___
6p
5d
4f
6s
5p
4d
5s
4p
3d
4s
3p
3s
2p
2s
1s
___ ___ ___
___ ___ ___ ___ ___
___ ___ ___ ___ ___ ___ ___
___
___ ___ ___
___ ___ ___ ___ ___
___
___ ___ ___
___ ___ ___ ___ ___
___
___ ___ ___
___
___ ___ ___
___
___
The number in front of the letter
represents the ENERGY LEVEL.
The letter represents the SUBLEVEL.
The superscript represents the number
of ELECTRONS in the sublevel.
1s2
http://intro.chem.okstate.edu/WorkshopFolder/Electronconfnew.html
ORBITAL DIAGRAMS
The boxes in the orbital diagram represent the ORBITALS.
TWO electrons with OPPOSITE spin can fit in each orbital.
The ARROWS represent the electrons. One arrow points
up and the other points down because the electrons in the
orbital have OPPOSITE spin.
Use the diagonal rule to determine the order of filling. Fill
the lowest energy sublevels first. Fill each degenerate
sublevel with single arrows before pairing. 4s sublevel is
lower in energy than the 3d so it is filled before the 3d.
C (Z = 6)




1s
2s
2p
3s
3p
4s
3d

S (Z = 16)


  

 

  

  
Mn (Z = 25)

1s

2s
2p
3s

3p



4s


3d
Half filled or filled sublevels are very STABLE.
Sometimes electrons do not follow the diagonal rule if the new
arrangement gives a FULL or HALF FULL sublevel.
Chromium (Z=24) is an exception to the diagonal rule because
of this. Draw the orbital diagram for chromium in the first set of
boxes below following the diagonal rule. In the second set of
boxes use the rule for more stable full of half-filled sublevels.


  



  

1s
2s
2p
3s
s
  





  





3p
4s

3d
http://www.wwnorton.com/chemistry/overview/ch3.htm
Shortcut
Move BACKWARD on the periodic table to the next NOBLE
GAS. Write the SYMBOL for the gas and enclose it in
brackets. Determine the OUTERMOST energy level by the
PERIOD. Move across the period assigning energy levels
and sublevels.
EX. The short hand Dubnium (105) would be found like this.
The next lower noble gas is RADON. Place its symbol in
brackets [Rn]. Move across the period assigning position to
the electrons. The next 2 electron would go into the 7s2. The
next 14 would go into the 5f14. The f sublevels are 2 energy
levels lower than their period. The next 3 would go into the
6d3. The d sublevels are 1 energy level lower than their
period.
The shorthand electron configuration for Db is
[Rn]7s2 5f14 6d3
Electron Dot Diagrams
1. Let the element's SYMBOL represent the nucleus and
inner energy level ELECTRONS.
2. Write the ELECTRON CONFIGURATION for the element
(or look at the PERIODIC table). Count the electrons in the
OUTER energy level.
3. Let the top, bottom, and two sides represent the
ORBITALS.
Place the appropriate number of DOTS in each "orbital".
Use ONE dot if the orbital has 1 electron and TWO dots if
the orbital has 2 electrons. Draw the electron dot diagrams
for:
Na(Z=11) 1s2 2s2 2p6 3s1
Na .
Mg(Z=12) 1s2 2s2 2p6 3s2
As(Z=33) 1s





2
2s
2
  
6
2
.Mg.
6
2
10
2p 3s 3p 4s 3d

  

4p
3
.
. As :
.
    
The periodic table can also be used to write electron dot
diagrams.
All the elements in the first column have ONE outer
electron(s). They have ONE dot(s) in their electron dot
diagram. All of the elements in column 2 have TWO outer
electrons. They have TWO dots in their electron dot
diagram. Most of the elements in column 18 have EIGHT
outer electrons except HELIUM which only has TWO outer
electrons.
Fill in the electron dot diagrams for each element on the
periodic table below
1
H
sH
e
Li
Be
B
C
N
O
F
N
e
3
Na
Mg
Al
Si
P
S
Cl
Ar
4
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
G
a
Ge
As
Se
Br
Kr
5
Rb
Sr
Y
Zr
N
b
M
o
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
6
Cs
Ba
Lu
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
7
Fr
Ra
Lr
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
2
ns
(n-1) d
np
La
Ce
Pr
N
d
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
N
o
(n-2) f
0
36
4 Sy 1
7
2
58