Download Ch. 4 Modern Chem Electrons

Chapter 4
Electrons
In Atoms
Properties of Light
 Electromagnetic
Radiation: a form of
energy that exhibits wavelike
behavior as it travels through space.
(ex. Visible light, x-rays, UV, IR,
radio)
 Electromagnetic Spectrum: all forms
of electromagnetic radiation together
make up the spectrum.
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Properties of Light
Characteristics of Waves
 There

are 4 main characteristics of
waves:
1) Amplitude: The height of the wave
measured from it’s origin to it’s peak.
When you increase the intensity, or
brightness of light, you are increasing
it’s amplitude.
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Characteristics of Waves
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Characteristics of Waves
 2)
Wavelength: the distance
between corresponding points on
adjacent waves. Wavelength is
designated by the Greek symbol
lambda.
Wavelength = λ
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Characteristics of Waves
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Characteristics of Waves
 3)
Frequency: the number of waves
that pass a given point in a specific
time. Measured in cycles per second
(cycle/second, or s-1) The SI unit for
this is Hertz. 1.0 Hz = 1.0 s-1
Frequency is designated by the Greek
symbol nu. Frequency = υ
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Characteristics of Waves
Characteristics of Waves
 4)
Speed: the speed of light is
constant. It is rounded to
8
3.00 x 10 m/s. The speed of
light is represented by the
letter c.
c = λυ
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The Photoelectric Effect
 Refers to the emission of electrons from a
metal when light shines on the metal.
 Quantum: the minimum amount of energy
that can be lost or gained by an atom
 Max Planck, a German physicist studied
quanta of light and found:
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The Photoelectric Effect
E = hυ
Where h is Plank’s
Constant and has a value
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of 6.6262 x 10 Js
(energy)
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The Photoelectric Effect
Einstein expanded upon this to propose that
light has a dual nature, acting as a wave
under some circumstances and a particle
under others.
Photon: a particle of electromagnetic
radiation having zero mass and carrying a
quantum of energy.
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Line Emission Spectrum
Ground State: the lowest energy state
of an atom
Excited State: any energy state that is
above ground state (always higher
energy than ground state).
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Line Emission Spectrum
Emission Line Spectrum: a graph that
indicates the degree to which a
substance emits radiant energy with
respect to wavelength.
Continuous Spectrum: the emission of
a continuous range of frequencies of
electromagnetic radiation.
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H Emission Spectrum
Emission Line Spectrum: a graph that
indicates the degree to which a
substance emits radiant energy with
respect to wavelenth.
Continuous Spectrum: the emission of
a continuous range of frequencies of
electromagnetic radiation.
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H Emission Spectrum
Emission Line Spectrum: a graph that
indicates the degree to which a
substance emits radiant energy with
respect to wavelenth.
Continuous Spectrum: the emission of
a continuous range of frequencies of
electromagnetic radiation.
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The Quantum Model
Heisenberg’s Uncertainty Principle: states
that it is impossible to determine
simultaneously both the position and
velocity of an electron or any other
particle.
Quantum Numbers: specify the
properties of atomic orbitals and the
properties of electrons in orbitals.
The Quantum Model
Principal Quantum Number (n): indicates
the main energy level occupied by the
electron. Values are positive integers only
(1, 2, 3, 4… with 1 being the lowest energy
level closest to the nucleus)
Angular Momentum Quantum Number (l):
indicates the shape of the orbital. Values
are 0, 1…n-1)
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The Quantum Model
l=0=s
l=1=p
l=2=d
l=3=f
The Quantum Model
Magnetic Quantum Number (m): indicates
the orientation of an orbital around the
nucleus. Values, including zero, are –l to
+l
l = 0 = s orbital has only one orientation (sphere)
l = 1 = p has three orientations
l = 2 = d has five orientations
l = 3 = f has seven orientations
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The Quantum Model
The Quantum Model
The Quantum Model
The Quantum Model
Spin Quantum Number: indicates the two
fundamental spin states of an electron in
an orbital. Values are +1/2 or -1/2.
Electron Configurations: the arrangement
of electrons in an atom.
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QUANTUM
NUMBERS
Values
Specifics
n
1,2,3,…..
l
n-1
0,1,2,3
0 = s orbital
1 = p orbital
2 = d orbital
3 = f orbital
m
-l, 0, +l
Orientations:
s has 1
p has 3
d has 5
f has 7
Ex: if “l” is 2, then “m”
can have values of
-2, -1, 0, +1, +2
s
-1/2, +1/2
Energy Level
*Period on Table
N/A
Orbital Filling Diagrams
There are 3 basic rules, named after the scientists
that discovered them, that govern the filling of
these orbitals with electrons…
1)
The Aufbau Principle: an electron
occupies the lowest energy orbital that
can receive it.
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Orbital Filling Diagrams
2) The Pauli Exclusion Principle: no two
electrons in the atom can have the same
set of four quantum numbers.
3) Hund’s Rule: All orbitals must be
occupied with a single electron before it
can be paired. Fill all orbitals with up
spin FIRST then fill with down spin
Orbital Filling Diagrams
- Each electron is designated as an arrow ()
- They can be Spin up ( ) or Spin down () - Each orbital is designated with a labeled
line: _____ or __ __ __ THESE LINES ARE THE # OF
1s
2p
ORIENTATIONS
-Multiple lines show multiple orbitals (1 for
s, 3 for p, 5 for d)
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Orbital Filling Diagrams
Orbitals fill going straight across each
period (row) on the periodic table, from
the lowest energy level up. (Aufbau).
Don’t forget, when they pair, they have
opposite spins (Pauli), but they won’t
pair until each available orbital has an
unpaired electron in it first (Hund)
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Orbital Filling Diagrams
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Orbital Filling Diagrams
-
-
-
The orbitals can hold a certain number of
electrons:
s=2
p=6
d=10
f=14
Orbital Filling Diagrams
Hydrogen, with one electron, would have an
orbital filling diagram of:

1s
Helium, with 2 electrons, would be:

1s
Now you’re at the end of the first period, start
again in the 2nd period with 2s…
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Orbital Filling Diagrams
Lithium:
 
1s 2s
Be:
 
1s 2s
B:
 
 _ _
1s 2s
2p
Orbital Filling Diagrams
Which of these would be correct for oxygen (with 8 e-):
O:
 
  ?
1s 2s
2p
OR
 
  
1s 2s
2p
The first one is correct, the second example violates the
Pauli Exclusion Principle.
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Practice
Write the orbital filling notation for the
following elements:
1) Be:_________________________
2) F:__________________________
3) Ar:_________________________
4) Cu:_________________________
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Electron Configurations
Now your ready to write electron configurations. These are simply the
orbital diagrams written out with superscripts:
Lithium:
 
1s 2s would be written out as 1s2 2s1
Be:
 
1s 2s would be written out as 1s2 2s2
B:
 
 _ _
1s 2s
2p
would be written out as 1s22s22p1
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Practice
Write the electron configuration for the
following elements:
5) Mg:________________________
6) N:_________________________
7) Cr:________________________
8) Cl:________________________
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Electron Configurations
Noble-Gas Configuration: refers to an outer main
energy level occupied by eight e

Once a subshell is complete at the end of a
period, you can write subsequent
configurations as having the core of the
previous noble gas with the additional valence
electrons.
Sodium Na would have a noble gas notation
of:
Na3s1
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Practice
Write the noble gas notation for the following
elements:
9) Ne:_________________________
10) Sb:_________________________
11) Y:__________________________
12) F:__________________________
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Ch. 4
The
End
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