The Wave Theory of Light and Electromagnetic Spectrum

undefined
 
Electronic Structure
and the Periodic Table
 
Unit 6 Honors Chemistry
 
Wave Theory of Light
James Clerk Maxwell
 
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Wave Diagram
 
Wave Vocab:
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1
 
Wave Vocab:
 
Amplitude 
– the distance from the origin to
the crest or the trough of a wave
Height (or intensity/brightness) of wave
Speed of light (c) 
– the rate at which all forms
of electromagnetic radiation travel through a
vacuum = 3.00 x10
8
 m/s
 
Wave Theory of Light
Comparing Waves
As Wavelength increases, frequency
_______________.
As Wavelength decreases, frequency
_______________.
 
Wave Equation
Wave Equation
 
One equation relates speed,
frequency and wavelength:
c = 
 
 
 
Example
 
The wavelength of the radiation which produces
the yellow color of sodium vapor light is 589.0 nm.
What is the frequency of this radiation?
 
c = 
 
 
The Electromagnetic Spectrum
 
Complete range of wavelengths and frequencies
Mostly invisible
 
What is Color?
 
TED Ed Video: What is color?
 
 
The Visible Spectrum
 
Continuous spectrum: components of white
light split into its colors, 
R
O
Y
 
G.
 
B
I
V
From 390 nm (violet) to 760 nm (red)
Can be split by a prism
 
How do we see color?
 
TED Ed Video: How we see color
 
Max Planck 
Particle Theory of Light
 
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P
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S
 
Equation
:
    
E = h
 
 
h =Plank’s constant= 6.626 x 10
-34 
J
·
s)
 
Example #1
 
(a) If the frequency of a ray of light is 5.09 x 10
14
Hz, calculate the energy, in joules, of a photon
emitted by an excited sodium atom.  (b) Calculate
the energy, in kilojoules, of a mole of excited
sodium atoms.
 
Example #2
 
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1
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m
?
 
Bohr Model of the Atom
 
When an electron absorbs a photon of energy,
the electron jumps from the ground state to an
excited state
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Line Spectra
Pattern of lines produced by light
emitted by excited atoms of an element
Unique for every element
Used to identify unknown elements
 
Explanation of Line Spectra
 
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B
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Energy of an electron is
quantized: 
can only
have specific values.
Energy is proportional
to energy level.
 
Explanation of Line Spectra
 
 
Electron will drop
from excited state
to ground state and
will emit energy as
a photon during the
fall.
 
Video: Atomic Emission Animation
 
Photoelectric Effect 
– Nobel Prize in
Physics 1921 to Einstein
 
Occurs when 
light 
strikes the surface of
a metal and 
electrons
 are ejected.
 
Practical uses
:
Automatic
door openers
 
Ted Ed Video: Is Light
Actually a Wave or Particle?
 
Conclusion…
 
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Light has a 
dual nature
!
 
Quantum Mechanics
 
Quantum mechanics:
 
atomic structure based on wave-
like properties of the electron
 
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Heisenberg Uncertainty
Principle
 
The exact location and speed of an electron
cannot be determined simultaneously (if we
try to observe it, we interfere with the
particle)
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Electrons exist in electron clouds and
not 
on specific rings or orbits like in
the Bohr model of the atom
 
Quantum Numbers
 
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They range from the most general locator to
the most specific
Analogy...
state = energy level, n
city = sublevel, l
address = orbital, m
l
house number = spin, m
s
 
1. Energy Level
Principal Quantum Number: n
 
Always a positive integer (1, 2, 3,…7)
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Similar to Bohr’s energy levels or shells
 
 
n = row number on periodic table for a given element
n in relation to the Periodic Table
 
Indicates shape of orbital
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Energy level 1 has only sublevel s
Energy level 2 has s and p
Energy level 3 has s, p, and d
Energy level 4-7 have s, p, d, and f
 
 
2. Sublevel
Angular Momentum Quantum Number: 
l
 
In order of increasing energy the
sublevels 
generally
 go:
 
s
 
<
 
p
 
<
 
d
 
<
 
f
 
HOWEVER, there
is some
overlapping of
sublevels at higher
energy levels
 
Ex.) 4s vs. 3d
 
3. Orbital
 
The most specific piece of information is about the
number and location of the electrons within the sublevel
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!
Orbitals
 
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No more than 2 e- assigned to an orbital
Orbitals grouped in s, p, d (and f) subshells
 
Shapes of Atomic Orbitals
 
 
s = spherical
p = peanut
d = dumbbell (clover)
f = flower
 
Capacities of levels, sublevels, and orbitals
 
 
Rules for how the electrons fill into the
electron cloud:
 
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Why are these incorrect?
Why are these incorrect?
Why are these incorrect?
Electron Configuration
 
Definition:
 
describes the distribution of
electrons among the various orbitals in
the atom
 
Represents the most
probable
 location of
the electron!
 
Electron Configurations
 
The system of numbers and letters that
designates the location of the electrons
3 major methods:
Full electron configurations
Abbreviated/Noble Gas configurations
Orbital diagram configurations
 
Full Electron Configuration
 
Example Notation:
1s
2
 2s
1  
  (Pronounced “
one-s-two, two-s-one
”)
A. What does the 
coefficient
 mean?
Principle energy level
B. What does the 
letter
 mean?
Type of sublevel – s, p, d, or f
C. What does the 
exponent
 mean?
# of electrons in that sublevel
 
Steps to Writing Full Electron
Configurations
 
1.
 
Determine the total number of electrons the atom
has (for neutral atoms it is equal to the atomic
number for the element).
  
Example:  F
   
atomic # =          # of p
+
 = # of e
-
 =
 
2.
 
Fill orbitals in order of increasing energy (see
Aufbau Chart).
 
3.
 
Make sure the total number of electrons in the
electron configuration equals the atomic number.
 
 
Aufbau Chart (Order of Energy Levels)
 
When writing electron
configurations:
d sublevels are n – 1
from the row they
appear in
 
f sublevels are n – 2
from the row they
appear in
 
In order of increasing energy the
sublevels 
generally
 go:
 
s
 
<
 
p
 
<
 
d
 
<
 
f
 
HOWEVER, there
is some
overlapping of
sublevels at higher
energy levels
 
Ex.) 4s vs. 3d
Writing Electron
Configurations
 
Nitrogen:
Helium:
Phosphorous:
Rhodium:
Bromine:
Cerium:
 
Abbreviated/Noble Gas Configuration
 
i.
 
Where are the noble gases on the periodic
table?
 
ii.
 
Why are the noble gases special?
 
iii.
 
 How can we use noble gases to shorten
regular electron configurations?
 
Abbreviated/Noble Gas Configuration
 
E
x
a
m
p
l
e
:
 
 
T
i
n
1
.
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e
l
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m
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i
s
.
 
2.
 
Start the configuration with the symbol for that
noble gas in brackets, followed by the rest of
the electron configuration.
 
Abbreviated/Noble Gas Configuration
 
P
r
a
c
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i
c
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!
 
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s
:
Sufur:
Rubidium:
Bismuth:
Zirconium:
Orbital Diagrams
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=
 
1
,
 
l
 
=
 
0
,
 
m
l
 
=
 
0
,
 
m
s
 
=
 
+
 
½
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,
 
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,
 
m
l
 
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0
,
 
m
s
 
=
 
-
 
½
 
Orbital Diagrams
 
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o
 
r
e
p
r
e
s
e
n
t
 
t
h
e
 
e
l
e
c
t
r
o
n
s
.
 
= orbital
 
sublevels
 
Orbital Diagrams
Don’t forget - 
orbitals have a capacity of two electrons
!!
Two electrons in the same orbital must have opposite
spins so draw the arrows 
pointing in opposite directions
.
    
Example:
  oxygen
  
1s
2
2s
2
2p
4
 
1s
 
2s
 
2p
 
Increasing Energy 
 
Drawing Orbital Diagrams
 
1.
 First, determine the electron configuration for the element.
2.
 Next draw boxes for each of the orbitals present in the electron
configuration.
Boxes should be drawn in order of increasing energy (see
the Aufbau chart).
3.
A
r
r
o
w
s
 
a
r
e
 
d
r
a
w
n
 
i
n
 
t
h
e
 
b
o
x
e
s
 
s
t
a
r
t
i
n
g
 
f
r
o
m
 
t
h
e
 
l
o
w
e
s
t
 
e
n
e
r
g
y
s
u
b
l
e
v
e
l
 
a
n
d
 
w
o
r
k
i
n
g
 
u
p
.
 
T
h
i
s
 
i
s
 
k
n
o
w
n
 
a
s
 
t
h
e
 
A
u
f
b
a
u
 
p
r
i
n
c
i
p
l
e
.
A
d
d
 
e
l
e
c
t
r
o
n
s
 
o
n
e
 
a
t
 
a
 
t
i
m
e
 
t
o
 
e
a
c
h
 
o
r
b
i
t
a
l
 
i
n
 
a
 
s
u
b
l
e
v
e
l
b
e
f
o
r
e
 
p
a
i
r
i
n
g
 
t
h
e
m
 
u
p
 
(
H
u
n
d
s
 
r
u
l
e
)
T
h
e
 
f
i
r
s
t
 
a
r
r
o
w
 
i
n
 
a
n
 
o
r
b
i
t
a
l
 
s
h
o
u
l
d
 
p
o
i
n
t
 
u
p
;
 
t
h
e
 
s
e
c
o
n
d
a
r
r
o
w
 
s
h
o
u
l
d
 
p
o
i
n
t
 
d
o
w
n
 
 
(
P
a
u
l
i
 
e
x
c
l
u
s
i
o
n
 
p
r
i
n
c
i
p
l
e
)
4.
 Double check your work to make sure the number of arrows in
your diagram is equal to the total number of electrons in the atom.
# of electrons = atomic number for an atom
Orbital Configurations for
Nitrogen
 
Full Electron Configuration:
 
 
Orbital Diagram:
Orbital Configurations for
Nickel
 
Full Electron Configuration:
 
 
Orbital Diagram:
 
Exceptions to the Filling Order Rule
(Cr, Cu)—these will not be on test!
 
Valence Electrons
 
Definition:
E
l
e
c
t
r
o
n
s
 
i
n
 
t
h
e
 
o
u
t
e
r
m
o
s
t
 
e
n
e
r
g
y
 
l
e
v
e
l
s
They determine the chemical properties of an
element!
 
***Write the noble gas configuration...the valence
electrons are the ones beyond the noble gas core
in the highest energy level
Valence Electrons and Core
Configuration (Shorthand)
 
What is the shorthand notation for S?
Configurations of Ions
 
Cations
: 
Formed when metals 
lose e
 in
highest principal energy level.
 
Example:
Configurations of Ions
 
Anions
: 
Formed when non-metals 
gain e
 to
complete the p sublevel
 
Transition Metals
 
Transition metals (and p block metals) 
lose e
from the highest principal energy level 
(
n
)
FIRST
, then lose their d electrons!
 
EOS
Zr    = 
[Kr] 5s
2
4d
2
Zr
+2
 = 
[Kr] 4d
2
undefined
 
Periodic Trends!
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
 
Electronegativity
Ability of an atom to pull e
-
 towards itself
Linus Pauling
: developed scale to demonstrate
different electronegativity strengths
 
Increases going 
up
 and to the 
right
Across a period
 
 more protons in nucleus =
more positive charge to pull electrons closer
Down a group
 
 more electrons to hold onto =
element can’t pull e
-
 
as closely
undefined
 
Electronegativity
Ability of an atom to pull e
-
 towards itself
Across a period
 
 more protons in nucleus =
more positive charge to pull electrons closer
Down a group
 
 more electrons to hold onto =
protons in nucleus can’t pull e
-
 
as closely
 
Periodic Properties & Trends
Periodic Properties & Trends
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
Atomic Radius
Distance between the nucleus and the
furthest electron in the valence shell
Increases going 
down
 and to the 
left
Down a group
 
 more energy shells = larger
radius
Across a period
 
 elements on the right can pull
e
-
 closer to the nucleus (more electronegative) =
smaller radius
 
*Remember*
LLLL  
 L
ower
, L
eft
, L
arge
, L
oose
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
Atomic Radius
Increases going
down
 and to the 
left
 
 
 
*Remember*
 
LLLL  
 
L
ower
, L
eft
,
 
L
arge
, L
oose
undefined
Memory Device
 
 
LLLL:  Lower Left, Larger Atoms
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
 
Ionic Radius
Radius of an atom when e
-
 are lost or
gained
 
different from atomic radius
 
Ionic Radius of 
Cations
Decreases when e- are removed
Ionic Radius of 
Anions
Increases when e- are added
undefined
Sizes of Ions
 
CATIONS are SMALLER than the
atoms from which they are formed.
 
Size decreases due to increasing he
electron/proton attraction.
L
i
,
1
5
2
 
p
m
3
e
 
a
n
d
 
3
p
undefined
Sizes of Ions
 
ANIONS are LARGER than the atoms
from which they are formed.
 
Size increases due to more electrons
in shell.
undefined
 
Trends in Ion Sizes
 
A
A
c
c
t
t
i
i
v
v
e
e
 
 
F
F
i
i
g
g
u
u
r
r
e
e
 
 
8
8
.
.
1
1
5
5
 
Trends in ion sizes are the same
Trends in ion sizes are the same
as atom sizes.
as atom sizes.
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
Ionization Energy
Energy required to remove an e- from the
ground state
1
st
 I.E. = removing 1 e
-
, easiest
2
nd
 I.E. = removing 2 e
-
, more difficult
3
rd
 I.E. = removing 3 e
-
, even more difficult
 
Ex.)  B  -->  B
+
  +  e-
 
       I.E. = 801 kJ/mol
Ex.)  B
+
  -->  B
+2
  +  e-
 
       I.E.2 = 2427 kJ/mol
Ex.)  B
+2
  -->  B
+3
  +  e-
 
       I.E.3 = 3660 kJ/mol
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
 
Ionization Energy
Increases going 
up
 and to the 
right
Down a group
 
 more e
-
 for the nucleus to
keep track of = easier to rip an e
-
 off
Across a period
 
 elements on the right
can hold electrons closer (more
electronegative) = harder to rip an e
-
 off
 
undefined
Memory Device
 
 
LLLL:  Lower Left,
   
Larger Atoms;
   
Looser electrons
undefined
 
Periodic Properties & Trends
Periodic Properties & Trends
 
 
Metallic Character
How “metal-like” an element is
Metals lose e
-
Most Metallic: Cs, Fr
Least: F, O
 
Increases going 
down
 and to the 
left
 
Think about where the metals & nonmetals are
located on the periodic table to help you remember!
undefined
Electron Affinity
 
Electron affinity 
is the energy involved
when an atom gains an electron to form an
anion.
Some elements 
GAIN
 electrons to form
anions
.
 
 
 
 A(g) +  e-  ---> A
-
(g)    E.A. = ∆E
undefined
 
Trends in Electron Affinity
 
Trend in a group:
Affinity for e
-
increases going
up a group
Trend in a series
or period:
Affinity for e
-
increases going
across a period to
the right
undefined
Electron Affinity
 
Note that the
trend for E.A.
is the SAME
as for I.E
.
!
 
A Summary of Periodic Trends
 
Remember LLLL!!
Slide Note

Dr. Mihelcic Honors Chemistry

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Explore the fascinating world of electromagnetic waves, visible light, and the wave theory through concepts such as wavelength, frequency, amplitude, and the speed of light. Understand how these elements are interconnected, and discover the diverse range of the electromagnetic spectrum. Dive into the relationship between wavelength and frequency and delve into practical examples to solidify your understanding.

  • Wave Theory
  • Electromagnetic Spectrum
  • Wavelength
  • Frequency
  • Speed of Light

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  1. Electronic Structure and the Periodic Table Unit 6 Honors Chemistry

  2. Wave Theory of Light James Clerk Maxwell Electromagnetic waves a form of energy that exhibits wavelike behavior as it travels through space Visible light a form of electromagnetic radiation that is perceivable to human beings and is seen in the colors of the rainbow ROY G. BIV

  3. Wave Diagram

  4. Wave Vocab: Crest the top of a wave Trough the bottom of a wave Wavelength ( crest to crest or trough to trough in a wave Units: m, nm (1 m = 109nm) lambda ) the distance from Frequency ( that pass a given point in a set amount of time (generally in 1 second) Units: Hertz (Hz), 1/s, or s-1 nu ) the number of wavelengths

  5. Wave Vocab: Amplitude the distance from the origin to the crest or the trough of a wave Height (or intensity/brightness) of wave Speed of light (c) the rate at which all forms of electromagnetic radiation travel through a vacuum = 3.00 x108m/s

  6. Wave Theory of Light

  7. Comparing Waves As Wavelength increases, frequency _______________. As Wavelength decreases, frequency _______________.

  8. Wave Equation One equation relates speed, frequency and wavelength: c =

  9. c = Example The wavelength of the radiation which produces the yellow color of sodium vapor light is 589.0 nm. What is the frequency of this radiation?

  10. The Electromagnetic Spectrum Complete range of wavelengths and frequencies Mostly invisible

  11. What is Color? TED Ed Video: What is color?

  12. The Visible Spectrum Continuous spectrum: components of white light split into its colors, ROY G. BIV From 390 nm (violet) to 760 nm (red) Can be split by a prism

  13. How do we see color? TED Ed Video: How we see color

  14. Max Planck Particle Theory of Light Light is generated as a stream of light particles called PHOTONS Equation: E = h h =Plank s constant= 6.626 x 10-34 J s)

  15. Example #1 (a) If the frequency of a ray of light is 5.09 x 1014 Hz, calculate the energy, in joules, of a photon emitted by an excited sodium atom. (b) Calculate the energy, in kilojoules, of a mole of excited sodium atoms.

  16. Example #2 What is the energy of a photon from the green portion of the rainbow if it has a wavelength of 4.90 x 10-7 m?

  17. Bohr Model of the Atom When an electron absorbs a photon of energy, the electron jumps from the ground state to an excited state Ground state lowest energy level an electron occupies Excited state temporary state when an electron is at a higher energy level

  18. Line Spectra Pattern of lines produced by light emitted by excited atoms of an element Unique for every element Used to identify unknown elements

  19. Niels Bohr Explanation of Line Spectra Niels Bohr Energy of an electron is quantized: can only have specific values. Energy is proportional to energy level.

  20. Explanation of Line Spectra Electron will drop from excited state to ground state and will emit energy as a photon during the fall. Video: Atomic Emission Animation

  21. Photoelectric Effect Nobel Prize in Physics 1921 to Einstein Occurs when light strikes the surface of a metal and electrons are ejected. Practical uses: Automatic door openers Ted Ed Video: Is Light Actually a Wave or Particle?

  22. Conclusion Light not only has wave properties but also has particle properties. These massless particles, called photons, are packets of energy. Light has a dual nature!

  23. Quantum Mechanics Quantum mechanics: atomic structure based on wave- like properties of the electron befo-schroedinger Schr dinger: wave equation that describes hydrogen atom

  24. Heisenberg Uncertainty Principle The exact location and speed of an electron cannot be determined simultaneously (if we try to observe it, we interfere with the particle) You can know either the location or the velocity but not both Werner Heisenberg Electrons exist in electron clouds and not on specific rings or orbits like in the Bohr model of the atom

  25. Quantum Numbers Quantum numbers a system of four numbers used to represent the most probable location of an electron in an atom They range from the most general locator to the most specific Analogy... state = energy level, n city = sublevel, l address = orbital, ml house number = spin, ms

  26. 1. Energy Level Principal Quantum Number: n Always a positive integer (1, 2, 3, 7) Indicates size of orbital, or how far electron is from nucleus Larger n value = larger orbital or farther distance from nucleus Similar to Bohr s energy levels or shells

  27. n in relation to the Periodic Table n = row number on periodic table for a given element n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7

  28. 2. Sublevel Angular Momentum Quantum Number: l Indicates shape of orbital Letters s, p, d, and f Energy level 1 has only sublevel s Energy level 2 has s and p Energy level 3 has s, p, and d Energy level 4-7 have s, p, d, and f

  29. In order of increasing energy the sublevels generally go: s < p < d < f HOWEVER, there is some overlapping of sublevels at higher energy levels Ex.) 4s vs. 3d

  30. 3. Orbital The most specific piece of information is about the number and location of the electrons within the sublevel The s sublevel has 1 orbital The p sublevel has 3 orbitals The d sublevel has 5 orbitals The f sublevel has 7 orbitals Orbital - region within a sublevel where an e- can be found (homes for e-) Every orbital can hold 2 electrons!

  31. Orbitals Orbital = electron containing area (houses for electrons) No more than 2 e- assigned to an orbital Orbitals grouped in s, p, d (and f) subshells

  32. Shapes of Atomic Orbitals s = spherical p = peanut d = dumbbell (clover) f = flower

  33. Capacities of levels, sublevels, and orbitals Maximum Number of Electrons in Energy Level Total Number of Orbitals Principal Energy level (n) Sublevels Present (s, p, d, or f) Number of Orbitals Present s p d f 1 2 3 4

  34. Rules for how the electrons fill into the electron cloud: Aufbau Principle: electrons fill from the lowest energy level to the highest (they don t skip around) Pauli ExclusionPrinciple: each orbital can hold a maximum of 2 electrons at a time (and they must have opposite spins) Hund s Rule: orbitals of equal energy in a sublevel must all have 1 electron before the electrons start pairing up

  35. Why are these incorrect?

  36. Why are these incorrect?

  37. Why are these incorrect?

  38. Electron Configuration Definition: describes the distribution of electrons among the various orbitals in the atom Represents the most probable location of the electron! EOS

  39. Electron Configurations The system of numbers and letters that designates the location of the electrons 3 major methods: Full electron configurations Abbreviated/Noble Gas configurations Orbital diagram configurations

  40. Full Electron Configuration Example Notation: 1s2 2s1 (Pronounced one-s-two, two-s-one ) A. What does the coefficient mean? Principle energy level B. What does the letter mean? Type of sublevel s, p, d, or f C. What does the exponent mean? # of electrons in that sublevel

  41. Steps to Writing Full Electron Configurations 1. Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element). Example: F atomic # = # of p+ = # of e- = 2. Fill orbitals in order of increasing energy (see Aufbau Chart). 3. Make sure the total number of electrons in the electron configuration equals the atomic number.

  42. Aufbau Chart (Order of Energy Levels) When writing electron configurations: d sublevels are n 1 from the row they appear in f sublevels are n 2 from the row they appear in

  43. In order of increasing energy the sublevels generally go: s < p < d < f HOWEVER, there is some overlapping of sublevels at higher energy levels Ex.) 4s vs. 3d

  44. Writing Electron Configurations Nitrogen: Helium: Phosphorous: Rhodium: Bromine: Cerium:

  45. Abbreviated/Noble Gas Configuration i. Where are the noble gases on the periodic table? ii. Why are the noble gases special? iii. How can we use noble gases to shorten regular electron configurations?

  46. Abbreviated/Noble Gas Configuration Example: Tin 1. Look at the periodic table and find the noble gas in the row above where the element is. 2. Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration.

  47. Abbreviated/Noble Gas Configuration Practice! Write Noble Gas Configurations for the following elements: Sufur: Rubidium: Bismuth: Zirconium:

  48. Orbital Diagrams Another way of writing configurations is called an orbital diagram. (also called orbital notation) ORBITAL BOX NOTATION for He, atomic number = 2 Arrows depict electron spin 2 1 s 1s One electron has n = 1, l = 0, ml = 0, ms = + Other electron has n = 1, l = 0, ml = 0, ms = -

  49. Orbital Diagrams Orbital diagrams use boxes (sometimes circles) to represent energy levels and orbitals. Arrows are used to represent the electrons. = orbital sublevels

  50. Orbital Diagrams Don t forget - orbitals have a capacity of two electrons!! Two electrons in the same orbital must have opposite spins so draw the arrows pointing in opposite directions. Example: oxygen 1s22s22p4 Increasing Energy 2p 2s 1s

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