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Arrangement of Electrons on Atoms
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Properties of light

Electromagnetic radiation
 Energy which travels through space as waves
 Speed is 3.0 x108 m/s
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Electromagnetic spectrum
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Gamma rays
X-rays
Ultra-violet rays
Visible lights
Infrared rays
Microwaves
Radio waves
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Wave length – the distance between
corresponding points on adjacent waves
Frequency (V)
The number of waves that pass a point per second
 Units are waves/second (often written as 1/s or s-1 )
 1 wave /second = 1 hertz (Hz)
 Speed = wavelength x frequency
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C=
V
C = 3.0 x 108 m/s
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Find the frequency of radiation if the
wavelength is 5.0 x 10-8 m
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Electrons are emitted when light strikes the
surface of certain metals
The light must be above a minimum frequency
for this to occur
Max Planck (1900) suggested that an object
emits energy in small specific amounts called
quanta
Quantum – the minimum energy that can be
lost or gained by an atom
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E = energy (J), H = 6.626 x 10-34Js ,
V = frequency
E=hv
Albert Einstein (1906) said that light has properties
of both waves and particles (photons)
Photons – a quantum of light energy
In order for the electron to be given off from the
metal, the electron must be struck by a photon
having enough energy to knock the electron loose
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Find the energy of radiation if the frequency is
6.15 x 1012 s-1
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Ground state – the lowest energy state of an
atom
Excited state - a higher energy state; when an
atom gains energy
When an atom returns to the ground state, it
gives off energy which may include visible
light
When the light emitted from hydrogen is
passed through a prism a series of lines of
lights of different wavelengths are seen

When an atom falls from the excited state to a
lower state a photon of radiation is emitted; the
energy of the photon is the difference in the
energy between the initial state and final state.
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The electron circles the nucleus in a definite
path (orbit)
Electron can move to a higher orbit by gaining
a certain amount of energy
When the electron drops down to a lower orbit,
a photon is emitted; the energy of a photon is
equal to the difference in energy between two
orbits
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Electrons as waves
Louis de Broglie (1924) proposed that electrons
have wave-like properties
Heisenberg Uncertainty Principal – it is
impossible to determine both the position and
path of the electron at the same time
Electrons are detected by their interaction with
photons
Because electrons are so small, an attempt to
locate an electron with a photon knocks the
electron off its course
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Erwin Schrodinger (1927) developed a
mathematical equation that treated electrons as
waves
Electrons do not travel in exact orbits; instead
they exist in certain regions of space called
orbitals
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A set of four quantum numbers describes the
properties of an atomic orbital.
Quantum Number
Describes
Principal (n)
Energy level
Angular momentum(ℓ) Shape of orbital
Magnetic (m)
orientation in space
Spin
spin of electron
Energy Level Sublevels #of obitals #of electron
1
s
1
2
2
s
1
2
p
3
6
3
s
1
2
p
3
6
d
5
10
4
s
1
2
p
3
6
d
5
10
f
7
14
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Electron configuration –the arrangement of the
electrons in an atom
Rules for Electron configuration
1. Aufbau Principle – an electron occupies the
lowest orbital possible
2. Pauli Exclusion Principal – an orbital may
contain 2 electrons which have opposite spin
3. Hund’s Rule – one electron enters each
orbital in a sublevel before a second electron
may enter any of them
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1s2 means there are 2 electrons in the 1s
sublevel
Use the periodic table as a guide when writing
electron configurations.
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H – 1s1
He – 1s2
Li – 1s22s1
O – 1s2 2s2 2p4
Na – 1s22s22p63s1
Ca – 1s22s22p63s23p64s2
Sc –1s22s22p63s23p64s23d1
Br – [Ar] 3d104s24p5
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1s
H( )
Li ( )
B( )
N( )
2s
( ) 2p
( ) ( )( )( )
( ) ( )( )( )
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Show the Symbol for the previous noble gas
and the electron config. which follows
Br [Ar] 4s23d104p5
Sb [Kr] 5s24d105p3
Exceptions to the pattern for the electron
arrangement
4s
3d
Cr [Ar] ( ) ( ) ( ) ( ) ( ) ( )
Cu [Ar] ( ) ( ) ( ) ( ) ( ) ( )

Sublevels which are filled or half filled are
more stabled then other configurations.