Chapter
Electrons in the Atom
Timeline of the Models*
?
460 BC
Democritus
“Atoms”
1800
John Dalton
“Billiard Ball”
1897
JJ
Thompson
“Plum
Pudding”
1911
Rutherford
“Planetary”
1912
Niels Bohr
“Bohr Model”
Today
Many
scientists
“Modern”
*There are many more models. These are the ones we’ll cover in class.
John Dalton (c. 1800 AD)

English chemist, John Dalton, performed
experiments with various chemicals and
showed matter seemed to consist of
“indivisible” particles (atoms).
 Though he didn’t know about atoms’
structure, he did know about the Law of
Conservation of Matter and based his theory
on this.
Dalton was an avid weather watcher and discoverer of color blindness
among other things.
Dalton’s Ideas
“Billiard Ball” or “Marbles”
Dalton's model says atoms are tiny,
indivisible, indestructible particles
 Believed each atom had a certain
mass, size, and chemical behavior
determined by what kind of elements
they make up.
 See next slide for the details….

John Dalton Theory --1800
(a.k.a “the marble guy”)

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
Atoms are the smallest
particles of nature-indivisible
and indestructible
All atoms of the same element
are identical
Atoms of different kinds can
combine to form compounds
Chemical reactions are atoms
recombining to form new
substances
J.J. Thomson (c. 1897)
In 1897, English physicist J.J.
Thompson discovered the electron
and proposed a model for the structure
of the atom.
 Using a CATHODE RAY TUBE,
Thomson discovered electrons have
a negative charge and thought that the
rest of matter must have a positive
charge to offset the negative electron.

His Experiment
JJ Thompson’s Model
“Plum Pudding”

Because the beam of light traveled
from to the positive end of the tube
he concluded that the light had a
negative charge
 Because the beam could push a
paddle wheel he concluded that
the particle had mass.
 Thompson's model says atoms
are positively charged spheres
with negatively charged electrons
randomly located throughout.
Side Trip…Alpha Particles!
 Around
this time scientists also
discovered alpha rays (particles),
which had a positive charge.
 Some physicists thought these
alpha particles were made up
of the positive parts of JJ
Thompson’s atom.
Ernest Rutherford (1911)

Rutherford as a student worked
under J.J. Thompson
supervision at the famous
Cavendish Laboratories.
 In 1911 Ernest Rutherford
bombarded atoms with alpha
rays to investigate the inside of
the atom.
 The results were, to say the
least, unexpected!
Gold Foil Experiment


Rutherford used
Radium as the source
of the alpha particles
and shot them at a thin
gold foil like aluminum
foil but made of gold
A fluorescent screen
sat behind the gold foil
on which he could
observe
the alpha particles’
impact.



When the particles bounced
back or were deflected
Rutherford reasoned that it
hit something massive and
positive. This mass became
know as the nucleus.
When the alpha particles
went straight through it hit
nothing.
This happened most often so
the atom is mostly empty
space.
Rutherford’s “Planetary Model



Rutherford’s model said the
negative electrons orbited a
positive center (NUCLEUS)like
our planets orbit the sun.
The nucleus contained most of the
mass of the atom
And the distance between the
positive center (nucleus) and the
electrons was huge-like a marble
in the center of a football field.
The atom was mostly empty
space!!!!
One little problem…

The theory of electricity
and magnetism predicted
that opposite charges attract
each other and the electrons
should gradually lose energy
and spiral inward toward the nucleus.
(BOOM! No more atom.)
Niels Bohr (1912)


In 1912 a Danish physicist, Niels Bohr
came up with a theory that said the
electrons do not spiral into the nucleus
and came up with some rules for what
does happen.
This was a pretty radical approach,
because for the first time rules had to fit
the observation regardless of how they
conflicted with the theories of the time!
(Aristotle would have been furious).
Previous experiments-White light
gives off all wavelengths of energy- all
colors.
The explanation
An electron absorbs
energy it jumps
farther away form the
nucleus
 As the electron falls
back closer to the
nucleus it gives off
the energy as colored
light.
How Light Relates to Electron Location

Bohr observed
that only certain
colors were given
off
 Therefore the
electron could
only orbit at
certain distances
from the nucleus
Bohr’s Rules

RULE 1: Electrons can orbit only at
certain allowed distances from the
nucleus (energy-levels).
 RULE 2: An atom absorbs energy when
an electron gets boosted from a lowenergy orbit to a high-energy orbit.
 Rule 3: Atoms radiate energy when an
electron jumps down from a higher-energy
orbit to a lower-energy orbit.
Bohr’s Model of the atom. Light can
excite electrons around atoms and
this gives rise to “quantum levels”.
Are We Done Yet?

Almost… Cliff Notes version is that
Niels Bohr came really close, and when
you add the works of Arnold
Sommerfeld, Wolfgang Pauli, Louis
de Broglie, Erwin Schrödinger,
Max Born, and Werner Heisenberg,
we arrive at today’s model…
Today’s Model!-Electron Cloud
Today's model
says electrons
are not confined to
fixed orbits.
 They occupy
volumes of
space outside
the nucleus.

Models of the Atom


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Dalton – indivisible sphere
Thomson Model – a ball of positive charge
containing a number of electrons “plum
pudding”
Rutherford Model – atom has a positively
charged nucleus
Bohr – electrons travel around the nucleus in
definite orbits (energy levels)
Quantum Mechanical Model – no definite
shape and no definite electron orbitals
Models of the Atom
Energy level – the region around the
nucleus where the electron is likely to be
moving (discrete levels, like stair-steps)
 Quantum – the amount of energy
required to move an electron from its
present energy level to the next higher
one

Models of the Atom


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Atomic orbital – cloud shaped regions where
electrons are thought to be located
S – orbital = Spherical
P – orbital = Peanut
D – orbital = Double peanut
F – orbital = Far too complex (Flower)
Nodes – regions where the probability of
finding an electron is very low (eg: No
electrons)
Table 5.1 Summary of Principle Energy Levels, Sublevels, and Orbitals
Principle
Energy Level
n=1
Number of
Sublevels
1
Types of
Orbitals
1s (1 orbital)
n=2
2
2s (1 orbital), 2p
(3 orbitals)
n=3
3
3s (1 orbital), 3p
(3 orbitals), 3d (5
orbitals)
n=4
4
4s (1 orbital),
4p (3 orbitals), 4d
(5 orbitals) 4f (7
orbitals)
Increasing energy
(increasing distance from nucleus)
Energy Level n
Maximum
number of
electrons
allowed
1
2
3
4
2
8
18
32
Electron Arrangement in Atoms
Electron Configurations – the way
electrons are arranged around the
nucleus of atoms
 Aufbau Principle – electrons enter
orbitals of lowest energy first
 Pauli exclusion Principle – An atomic
orbital may describe at most two
electrons

Electron Arrangement in Atoms
Hund’s rule – When electrons occupy
orbitals of equal energy, one electron
enters each orbital until all orbitals
contain one electron with parallel spins
 Once all orbitals of equal energy have
one electron with parallel spin, the next
electron to enter the orbital has the
opposite spin.

Electron Arrangement in Atoms

Cr =

Cu =
SEE OVERHEAD
Physics and the Quantum
Mechanical Model

Electromagnetic radiation – includes radio
waves, microwaves, visible light, infrared
light, ultraviolet light, X-rays, and gamma
rays
Physics and the Quantum
Mechanical Model
Amplitude – the height of the wave from the
origin to the crest
 Wavelength – the distance between the crests
 Frequency – the number of wave cycles to
pass a given point per unit of time

Crest
Physics and the Quantum
Mechanical Model
Hertz – the metric unit of cycles per
second
 Spectrum – when sunlight passes
through a prism, the light separates into
a spectrum of colors
 Atomic emission spectrum – passing
light emitted by an element through a
prism gives a spectrum of the element

Physics and the Quantum
Mechanical Model

c=‫ג‬xv

c = speed of light

v = frequency

‫ = ג‬wavelength
Physics and the Quantum
Mechanical Model
Photons – light particles
 Photoelectric effect – electrons called
photoelectrons are ejected by metals
when light shines on them.


Li, Na, K, Cs, and Rb are the best metals to
demonstrate the effect.
Physics and the Quantum
Mechanical Model
Ground State – the lowest energy level
(n = 1)
 Energy levels above the ground state are
denoted N = 2,3,4,5,6,ect…
 Electrons absorb energy to move up an
energy level and give off energy to move
down an energy level
