Transcript Slide 1

Chapter 6
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Arrangement of Electrons in Atoms
0r….
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Matter waves and waves that don’t matter
The nature of light
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Dual nature of light
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Wave characteristics   Particle characteristics
Wave nature of light
Electromagnetic radiation
Electromagnetic (EM) radiation
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Form of E w/ wavelength (l) behavior
Speed = 3.0 x 1010 cm/s (speed of light)
Wavelength (l)  distance between pts.
on a wave
Frequency (n)
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# of waves that pass a given pt. in a
specific time
frequency
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C = ln
Therefore, as l decreases, n increases
C = speed of light, (186,000 miles/s, or
299,792,458 m/ s)
Continuous spectrum
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All ls in a given range included
Electromagnetic (EM) spectrum
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All EM radiation
Particle nature of light
Photoelectric effect
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Emission of e- by certain metals when light
shines on them
Max Planck (1900)
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When a hot object loses E, it is lost in sm.
Specific amts. Called quanta
Quantum
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Finite quantity of E that can be gained or
lost by an atom
Photon
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Individual quantum of light
Albert Einstein (1905)
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Higher n = higher E
Absorb. Of photons of a specific E explains
photoelectric effect
Dual (wave-particle) nature of light
Important formulas
E = hn
h (Plank’s constant) = 6.626 x 10 -34 J . S
n (frequency)
c = ln
c (speed of light) = 186,000 miles/s, or
299,792,458 m/ s
Hydrogen atom spectrum
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Pass high voltage through H2 gas  gas
glows pass light through prism bright
line spectrum
Bright line spectrum
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Due to e-s boosted to high E state (excited
state), then dropping to the ground state
Lines represent E given off when e-s drop
to ground state
Hydrogen spectrum
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E of photon = difference between ground
and excited state
spectroscope
Flame test
Bohr Model of the atom (1913)
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The Hydrogen e- can circle the H nucleus
only in certain orbits (like rungs of a
ladder)
Definite orbits occupied by
electron particles
Worked w/ H atom only
According to this theory
an electron moving
between orbits would
disappear from one
and reappear instantaneously in another
without visiting the space between 
“Quantum leap”
“An electron doesn’t exist until it is
observed”
“Until it is observed an electron must be
regarded as being at once everywhere and
nowhere”
Dennis Overbye
Schrödinger Model (1926)
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Wave properties of
atoms
Worked w/ all atoms
e- in orbitals 
e- clouds
Can not pinpoint
location of e- and
path at a given
instant  immutable
property of the
universe
Quantum numbers
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“Electron address”
Location of e-s in the atom
Quantum number 1
“Pennsylvania”
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Principle quantum number (main energy
level)
n= 1,2,3……7
Quantum number 2
“Hollidaysburg”
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Orbital quantum number (shape of orbital)
s,p,d,f
Quantum number 3
“N. Montgomery St.”
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Magnetic quantum number (orientation of
orbital about the nucleus)
Quantum number 4
“1510”
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Spin Quantum number (two possible
states of electron)
+1/2 or -1/2
Arrangement of electrons
Arrangement of electrons
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Main energy levels: 1,2,3…..
Sublevels: s,p,d,f
Orbitals
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Each
Each
Each
Each
s has 1
p has 3
d has 5
f has 7
Each orbital can hold a max of 2 e-
Orbital notation
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Unoccupied orbital ___
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Orbital with 1 e-
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Orbital with 2 e-
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e.g. H
#
1s
# or $
#$
He
#$
1s
Orbital notation
Electron configuration notation
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Uses superscripts instead of lines
e.g. H 1s1
or He 1s2
Electron dot notation
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Uses only e- in highest (outermost) main
energy levels
e.g.
.Na
.
.He
Aufbau (building up) principle
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Electrons occupy lowest energy orbital
that will receive them, e.g. hydrogen’s
electron goes into the 1s orbital
Hund’s rule
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Orbitals of equal E are each occupied by
one e- before 2nd e- is added, all e- in
singly occupied orbitals must have same
spin
Pauli exclusion principle
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No two e- in same atom have the same set
of four quantum numbers
Electron fill chart
Shorthand notation
exceptions
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e.g. copper
[Ar] 4s1 3d10