Arrangement of
Electrons in Atoms
The Development of a New
Atomic Model
Light
 Before
1900, scientists thought that
light behaved only as wave
 discovered
that also has particle-like
characteristics
Light is electromagnetic
radiation.
 electromagnetic


radiation:
form of energy that acts as a wave as it
travels
includes: gamma, X rays, ultraviolet (UV),
visible and infrared (IR) light, microwaves,
and radio waves
 All
forms are combined to form
electromagnetic spectrum
Light as a Wave
Wave Properties:
 all
form of EM radiation travel at a
speed (c) of 3.0 x 108 m/s in a vacuum
 wavelength:
(λ) distance between points
on adjacent waves; in nm (109nm = 1m)
 frequency:
(f) number of waves that
passes a point in a second, in
waves/second
Inversely proportional!
c = λf
Waves
long wavelength l
Amplitude
Low
frequency
short wavelength l
Amplitude
High
frequency
Quantum Theory

Max Planck
(1900)
German physicist



Max Planck
Observed - emission
of light from hot
objects
Concluded –energy Planck’s constant (h) =
6.626 x 10-34 J*s
is emitted in small,
specific amounts
(quanta)
Quantum - minimum
amount of energy
change
E = hf
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Photoelectric Effect
 Albert

Einstein (1905)
Observed - photoelectric effect
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Albert Einstein
Photoelectric Effect
Electrons are emitted
No electrons are emitted
Bright
red light
or
Dim
blue light
or
infrared rays
ultraviolet rays
Metal plate
Metal plate
Photoelectric Effect
 when
light is shone on a piece of
metal, electrons can be emitted
 no
electrons were emitted if the light’s
frequency was below a certain value
 scientists
could not explain this with
their classical theories of light
Solar Calculator
Solar Panel
Photoelectric Effect
 Einstein
in 1905
added on to Planck’s theory
 suggested
that light can be viewed as
stream of particles
 photon-
particle of EM radiation
having no mass and carrying one
quantum of energy
Photoelectric Effect
 EM
radiation can only be absorbed by
matter in whole numbers of photons
 when metal is hit by light, an electron
must absorb a certain minimum
amount of energy to knock the
electron loose
 this minimum energy is created by a
minimum frequency
 since electrons in different metal
atoms are bound more or less tightly,
then they require more or less energy
H Line-Emission Spectrum

Why had hydrogen atoms only given off
specific frequencies of light?
current Quantum Theory attempts to
explain this using a new theory of atom
H Line-Emission Spectrum

ground state- lowest energy state of an atom

excited state- when an atom has higher
potential energy than it has at ground state
(excited by heat or electricity)

line-emission spectrum- pattern of
wavelengths of light created when visible
light from excited atoms is shined through a
prism
Line-Emission Spectrum

pattern of wavelengths of light created
when visible light from excited atoms is
shined through a prism
excited state
Wavelength (nm)
410 nm
486 nm
434 nm
Slits
ENERGY IN
Prism
PHOTON OUT
ground state
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
656 nm
H Line-Emission Spectrum
 scientists
using classical theory expected
atoms to be excited by whatever energy
they absorbed
 continuous spectrum- emission of
continuous range of frequencies of EM
radiation
H Line-Emission Spectrum
 when
an excited atom falls back to
ground state, it emits photon of
radiation
 the
photon is equal to the difference
in energy of the original and final
states of atom
 since
only certain frequencies are
emitted, the differences between the
states must be constant
Bohr Model
 created
by Niels Bohr
(Danish physicist)
in 1913
 linked atom’s electron with emission
spectrum
 electron
can circle nucleus in certain
paths, in which it has a certain amount
of energy
Bohr Model
 Can
gain energy by
moving to a higher rung
on ladder
 Can lose energy by
moving to lower rung
on ladder
 Cannot gain or lose
while on same rung of
ladder
Bohr Model
a photon is
released that
has an energy
equal to the
difference
between the
initial and final
energy orbits
Bohr Model
6
5
4
Energy
3
2
1
nucleus
of photon
depends on the
difference in energy
levels
Bohr’s
calculated
energies matched
the IR, visible, and
UV lines for the H
atom
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

http://www.teacher
sdomain.org/asset/
phy03_vid_quantu
m/

http://www.mhhe.c
om/physsci/chemis
try/essentialchemis
try/flash/linesp16.s
wf
Other Elements
 Each
element has a unique bright-line
emission spectrum.
i.e. “Atomic Fingerprint”
Helium
Bohr’s calculations only worked for
hydrogen! 
Didn’t explain chemical behavior of atoms.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Electrons as Waves
 In
1924, Louis de Broglie
(French scientist)
 suggested the way quantized
electrons orbit the nucleus is similar to
behavior of wave
 electrons can be seen as waves confined
to the space around a nucleus
 waves could only be certain frequencies
since electrons can only have certain
amounts of energy
Uncertainty Principle
 In
1927 by Werner Heisenberg
(German theoretical physicist)
 electrons can only be detected by their
interaction with photons
 any attempt to locate a specific
electron with a photon knocks the
electron off course
 Heisenberg Uncertainty Principle- it is
impossible to know both the position
and velocity of an electron
Electrons as Waves
c
v
c  lv
l
E  hv
h
l
mv
hc
l
 mc
E
hc
l
2
E  mc
shows that anything with both mass and velocity
has a corresponding wavelength
2
Schrödinger Wave Equation
 In
1926, Erwin Schrödinger
(Austrian physicist)
 his equation proved that
electron energies are quantized
 only waves of specific energies
provided solutions to his equation
 solutions to his equation are called
wave functions
Ψ 1s 

1 Z 3/2 σ
π a0
e
Teacher, may I be excused?
My brain is full!