Electrons, Energy, & the Electromagnetic
Spectrum
Simplified, 2D Bohr Model:
Orbits,
paths, shells, rings = ENERGY LEVELS
As
the name “energy level” implies, there is a
specific
amount of energy associated with each
energy
level.
Electrons
are lazy – they will occupy the
location
that requires the least amount of
energy
Lowest
energy state = GROUND STATE
For hydrogen in the ground state…
~
Electron is located in energy level 1 (E 1) because E1 is closest to the
nucleus. Nucleus is positivelycharged & attracts negativelycharged e. Therefore, low amount of energy Nucleus needed to be in E1.
Electron
If energy (from an outside
source) is added to the atom…
Electron
jumps to a higher energy level because
energy
is absorbed. (Electron is now in the
EXCITED
STATE.)
As
soon as the electron jumps to a higher level,
Add
energy it
immediately falls back to a lower energy level.
When
the electron falls, energy is released.
energy
The
energy is released in the form of
released electromagnetic
(EM) radiation.
The
type or form of EM radiation released
depends
on the difference in energy between
the
energy levels.

Here’s how the type/form of EM radiation can be determined…
The amount of energy released
when an electron falls from a higher to a lower energy level is directly
proportional to its frequency. The
calculation follows the equation: E
= h ^{.} ν
E = Energy (unit is J)
h = Planck’s Constant (6.626 x 10^{34} J^{.}s)
ν = frequency (unit is Hz or ^{1}/_{second})
EXAMPLE 1: A particle of EM radiation has an energy of
1.15 x 10^{16} J. What is its
frequency?
1.15
x 10^{16} J = 6.626 x 10^{34} J^{.}s ^{.}
ν
ν = 1.74 x 10^{17}
Hz
The type of electromagnetic
radiation can be determined if one knows the wavelength. The wavelength is inversely proportional to
the frequency. The calculation follows
the equation: c = λ ^{.}
ν
c = speed of light (3.00 x 10^{8}
m/s)
λ = wavelength (unit is m)
ν = frequency (unit is Hz or ^{1}/_{s})
EXAMPLE 2: What is the wavelength of the same particle
from EXAMPLE 1?
3.00
x 10^{8} m/s = λ ^{.} 1.74 x 10^{17} Hz
λ
= 1.72 x 10^{9} m
EXAMPLE 3: What type of electromagnetic radiation is the
particle from EXAMPLE 1?
answer for wavelength is 10^{9} so use chart below to determine…
xrays or ultraviolet
(either one is acceptable)
PROBLEMS FOR YOU TO TRY ON YOUR OWN…
1.) A particle of EM radiation
has a frequency of 5.76 x 10^{14} Hz.
(A) How much energy does this
particle have?
(B) What is the wavelength of
this particle?
(C)
What specific type of electromagnetic radiation does this particle represent?
2.) A particle of electromagnetic
radiation has 2.39 x 10^{13} Joules of energy.
(A) What is the wavelength of
this particle?
(B)
What type of electromagnetic radiation does this particle represent?

Light Calculations Notes:
* Frequency and wavelength are ___________________ proportional
* Energy and frequency are ____________________ proportional
Light as a Particle
Notes:
* Object emits energy in small, specific amounts (called ________________)
* _________________: particle of EM radiation carrying a quantum of energy
Quantum Theory
Notes:
* When atom falls from excited state to ground state, _______________
______________________________
____________________________________________________________
Bohr model of the
hydrogen atom
* drawback of Bohr’s model =
EMISSION & ABSORPTION SPECTRA
According
to the Bohr atomic model, electrons orbit the nucleus within specific energy
levels. These levels are defined by unique amounts of energy.
Electrons possessing the lowest energy are found in the levels closest to the
nucleus. Electrons with higher energy are located in progressively more
distant energy levels.
If an
electron absorbs sufficient energy to bridge the "gap" between energy
levels, the electron may jump to a higher level. Since this change
results in a vacant lower orbital, this configuration is unstable. The
"excited" electron releases its newly acquired energy as it falls
back to its initial or ground state. Often, the excited electrons
acquire sufficient energy to make several energy level transitions. When
these electrons return to the ground state, several distinct energy emissions
occur. The energy that the electrons absorb is often of a thermal or
electrical nature, and the energy that an electron emits when returning to the
ground state is called electromagnetic radiation.
In 1900,
Max Planck studied visible emissions from hot glowing solids. He
proposed that light was emitted in "packets" of energy called quanta,
and that the energy of each packet was proportional to the frequency of the
light wave. According to Einstein and Planck, the energy of the packet
could be expressed as the product of the frequency (n) of emitted light and Planck's
constant (h). The equation is written
as E = hn
If white
light passes through a prism or diffraction grating, its component wavelengths
are bent at different angles. This process produces a rainbow of distinct
colors known as a continuous spectrum. If, however, the light
emitted from hot gases or energized ions is viewed in a similar manner,
isolated bands of color are observed. These bands form characteristic
patterns  unique to each element. They are known as bright line
spectra or emission spectra.
By
analyzing the emission spectrum of hydrogen gas, Bohr was able to calculate the
energy content of the major electron levels. Although the electron
structure as suggested by his planetary model has been modified according to
modern quantum theory, his description and analysis of spectral emission lines
are still valid.
In addition
to the fundamental role of spectroscopy played in the development of today's
atomic model, this technique can also be used in the identification of
elements. Since the atoms of each element contain unique arrangements of
electrons, emission lines can be used as spectral fingerprints. Even
without a spectroscope, this type of identification is possible since the major
spectral lines will alter the color.
ELECTRON
ARRANGEMENT
Heisenberg
Uncertainty Principle:

GENERAL LOCATION  ____________________

V
____________________

V
____________________

V
SPECIFIC
LOCATION  ____________________
ENERGY
LEVELS  divisions of the electron
cloud
 numbered
consecutively from closest to farthest away from nucleus
.
Electrons will occupy the location with the lowest
amount of energy.
SUBLEVELS  divisions of energy levels
 designated by
letters (s, p, d, f)
 number of
sublevels in an energy level = # of the energy level

ORBITALS  divisions of sublevels
 number of orbitals in an
energy level =
 “s” sublevel has 1 orbital; “p”
sublevel has 3 orbitals; “d”
sublevel has 5 orbitals;
“f” sublevel has 7 orbitals
HOW
MANY ELECTRONS CAN AN ORBITAL HOLD?
HOW
MANY ELECTRONS CAN EACH SUBLEVEL HOLD?
“s” = ___ e “p” = ___ e “d” = ___ e “f” = ___ e
HOW
MANY ELECTRONS CAN AN ENERGY LEVEL HOLD?
AUFBAU
PRINCIPLE –
What
is the order in which the sublevels fill with electrons? Use the DIAGONAL RULE.
1s
2s 2p
3s 3p
3d
4s 4p
4d 4f
5s 5p
5d 5f
6s 6p
6d 6f
7s 7p
More
Electron Arrangement!
Examples:
Se 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
Sn 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{2}
Hg 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}
HOEL
(Highest Occupied Energy Level): energy
level furthest from the nucleus that contains at least one electron
How to determine this using electron
configuration?
~
largest nonexponent number
Se 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4} HOEL = 4
Sn 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{2} HOEL
= 5
Hg 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}
HOEL = 6
Valence
Electrons: electrons in the HOEL
How to determine this using electron
configuration?
~
add up exponents of terms in HOEL
Se 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2}
3d^{10} 4p^{4}
HOEL
= 4 Valence electrons = 2 + 4 =
6
Sn 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2} 4d^{10} 5p^{2}
HOEL
= 5 Valence electrons = 2 + 2 =
4
Hg 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{6} 6s^{2}
4f^{14} 5d^{10}
HOEL
= 6 Valence electrons = 2
Noble Gas
Configuration: shortcut for electron
configuration
How is it written?
~
[ symbol for noble has closest to element with lower atomic # ]
~
[after brackets] next number is the period that the element is located in
~
after that number, write “s”
~
continue electron configuration in diagonal rule order until appropriate # of
electrons is reached
*NOTE: ending of electron configuration and noble
gas configuration should be the same*
Se 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
[Ar] 4s^{2} 3d^{10}
4p^{4}
^{18 20 30 34}
Sn 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{2}
[Kr] 5s^{2} 4d^{10} 5p^{2}
^{ 36 38 48 50 }
Hg 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6} 5s^{2}
4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}
[Xe]
6s^{2} 4f^{14} 5d^{10}
^{ 54 56 70 80}
Orbital
Notation: drawing of how electrons are
arranged in orbitals; will only need to do this for the HOEL
*NOTE: ___
= orbital
or =
electrons
Se 1s^{2} 2s^{2} 2p^{6}
3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
Sn 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{2}
Hg 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}

Dot Diagrams: symbol represents nucleus and nonvalence
(“innershell”) electrons; dots around
symbol represent valence electrons
^{ }
^{}^{}
^{ }
^{ }
^{ }
^{ }
^{ }
^{ }
^{}^{ }
QUANTUM NUMBERS
~
describe one specific electron
~ 1st
quantum number = PRINCIPAL QUANTUM NUMBER
~ abbreviated
"n"
~ tells the
energy level the electron is located in
~ n = number of
the energy level
~ 1st energy
level: n = 1, 4th energy level: n = 4, etc.
~ 2nd quantum number
= ANGULAR MOMENTUM QUANTUM NUMBER
~ abbreviated " l "
~ tells the sublevel the
electron is located in
~ tells shape of orbital
~ "s" sublevel: l = 0, "p" sublevel: l = 1, "d"
sublevel: l = 2, "f"
sublevel: l = 3
~ 3rd quantum number
= MAGNETIC QUANTUM NUMBER
~ abbreviated "m"
~ tells which orbital the
electron is in
~ tells orientation of orbital
around nucleus
~ m =  l .. + l
~ abbreviated " s "
~ tells which electron is being
described
~ tells which way electron is
spinning
~ s = ^{1}/_{2} or +^{1}/_{2}
Pauli Exclusion Principle:
EXAMPLE QUESTIONS:
1.) What are the 4 quantum numbers for
the following electron?
h
3 p
2.) If the electron in question 1 was the last
electron added, what element would it be?
3.) Draw in the electron (and the orbital
notation) for the electron with the following quantum numbers.
n
= 3 l = 2 m = 1 s =  ½
4.) How many electrons in an atom can have the
quantum numbers n = 3 and l = 1?
5.) What are the four quantum numbers
for the electron circled in the diagram below?
n
= l
= m = s =