DEVELOPMENT OF A NEW ATOMIC MODEL
I.
Behavior of electrons related to absorption and emission of light
II. Properties of light
measured in ____ - abbreviated ____)
measured in ____ - abbreviated ____)
e.
frequency and wavelength are
___________________ proportional
III. Photoelectric effect
IV. Hydrogen atom line emission spectrum
specific frequencies of light (corresponds
to different colors)
V. Bohr model of the hydrogen atom
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
GENERAL LOCATION
-------------------- ____________________
|
V
____________________
|
V
____________________
|
V
SPECIFIC
LOCATION -------------------- ____________________
ENERGY
LEVELS -
- numbered from
closest to farthest away from nucleus
nucleus .
E 1 E 2 E 3 E 4
Electrons will occupy the location with the
________________________ amount of energy.
SUBLEVELS -
- designated by letters
- number of sublevels in an energy level =
nucleus .
E 1 E 2 E 3 E 4
ORBITALS -
-
number of orbitals in an
energy level =
- ___ sublevel has 1 orbital; ___ sublevel has 3 orbitals;
___ sublevel has 5 orbitals; ___ sublevel has 7 orbitals
HOW MANY
ELECTRONS CAN AN ORBITAL HOLD?
HOW MANY
ELECTRONS CAN EACH SUBLEVEL HOLD?
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 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p4
Sn 1s2
2s2 2p6 3s2 3p6 4s2 3d10
4p6 5s2 4d10 5p2
Hg 1s2 2s2 2p6
3s2 3p6 4s2 3d10 4p6 5s2
4d10 5p6 6s2 4f14 5d10
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 non-exponent number
Se 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p4 HOEL = 4
Sn 1s2
2s2 2p6 3s2 3p6 4s2 3d10
4p6 5s2 4d10 5p2 HOEL = 5
Hg 1s2 2s2 2p6
3s2 3p6 4s2 3d10 4p6 5s2
4d10 5p6 6s2 4f14 5d10
HOEL = 6
Valence Electrons: electrons in the HOEL
How to determine this using electron configuration?
~ add
up exponents of terms in HOEL
Se 1s2 2s2 2p6 3s2
3p6 4s2
3d10 4p4
HOEL
= 4 Valence electrons = 2 + 4 =
6
Sn 1s2
2s2 2p6 3s2 3p6 4s2 3d10
4p6 5s2
4d10 5p2
HOEL
= 5 Valence electrons = 2 + 2 =
4
Hg 1s2 2s2 2p6
3s2 3p6 4s2 3d10 4p6 5s2
4d10 5p6 6s2
4f14 5d10
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 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p4
[Ar]
4s2 3d10 4p4
18 20
30 34
Sn 1s2
2s2 2p6 3s2 3p6 4s2 3d10
4p6 5s2 4d10 5p2
[Kr] 5s2 4d10 5p2
36
38 48 50
Hg 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6 5s2 4d10
5p6 6s2 4f14 5d10
[Xe] 6s2 4f14 5d10
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 h or i = electrons
Se 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p4 hi hi
h h
4s 4p
Sn 1s2
2s2 2p6 3s2 3p6 4s2 3d10
4p6 5s2 4d10 5p2 hi h h
__
5s 5p
Hg 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6 5s2 4d10
5p6 6s2 4f14 5d10 hi
6s
Dot Diagrams: symbol represents nucleus and non-valence
(“inner-shell”) electrons;
dots around symbol represent valence electrons