Atomic Models

⚫ This model of the

atom may look
familiar to you. This is
the Bohr model. In
this model, the
nucleus is orbited by
electrons, which are
in different energy
levels.
⚫ A model uses familiar ideas to
explain unfamiliar facts
observed in nature.
⚫ A model can be changed as
new information is collected.
⚫ The atomic
model has
changed
throughout the
centuries,
starting in 400
BC, when it
looked like a
billiard ball →
 Who are these men?
In this lesson, we’ll learn
about the men whose quests
for knowledge about the
fundamental nature of the
universe helped define our
views.
 Democritus
400 BC

⚫ This is the Greek

philosopher Democritus
who began the search for
a description of matter
more than 2400 years
ago.
⚫ He asked: Could
matter be divided into
smaller and smaller
pieces forever, or was
there a limit to the
number of times a
piece of matter could
be divided?
Atomos

⚫ His theory: Matter could

not be divided into
smaller and smaller
pieces forever, eventually
the smallest possible
piece would be obtained.
⚫ This piece would be
indivisible.
⚫ He named the smallest
piece of matter “atomos,”
meaning “not to be cut.”
Atomos

▪ To Democritus, atoms
were small, hard
particles that were all
made of the same
material but were
different shapes and
sizes.
▪ Atoms were infinite in
number, always
moving and capable
of joining together.
 This theory was ignored and
forgotten for more than 2000
years!
 Dalton’s Model

⚫ In the early 1800s,

the English
Chemist John
Dalton performed a
number of
experiments that
eventually led to
the acceptance of
the idea of atoms.
Dalton’s Theory
⚫ He deduced that all
elements are composed of
atoms. Atoms are
indivisible and
indestructible particles.
⚫ Atoms of the same element
are exactly alike.
⚫ Atoms of different elements
are different.
⚫ Compounds are formed by
the joining of atoms of two
or more elements.
 .

⚫ Thistheory
became one
of the
foundations
of modern
chemistry.
Thomson’s Plum Pudding
Model
⚫ In1897, the
English scientist
J.J. Thomson
provided the first
hint that an atom
is made of even
smaller particles.
 Thomson Model
⚫ He proposed a
model of the atom
that is sometimes
called the “Plum
Pudding” model.
⚫ Atoms were made
from a positively
charged substance
with negatively
charged electrons
scattered about,
like raisins in a
pudding.
 Thomson Model
⚫ Thomson studied
the passage of
an electric
current through a
gas.
⚫ As the current
passed through
the gas, it gave
off rays of
negatively
charged
particles.
 Thomson Model
Where did
they come
⚫ This
surprised from?

Thomson,
because the
atoms of the gas
were uncharged.
Where had the
negative charges
come from?
Thomson concluded that the
negative charges came from within
the atom.

A particle smaller than an atom had

to exist.

The atom was divisible!

Thomson called the negatively
charged “corpuscles,” today known
as electrons.

Since the gas was known to be

neutral, having no charge, he
reasoned that there must be
positively charged particles in the
atom.

But he could never find them.

 Rutherford’s Gold Foil
Experiment
⚫ In 1908, the
English physicist
Ernest Rutherford
was hard at work
on an experiment
that seemed to
have little to do
with unraveling the
mysteries of the
atomic structure.
⚫ Rutherford’s experiment Involved
firing a stream of tiny positively
charged particles at a thin sheet of
gold foil (2000 atoms thick)
⚫ Most of the positively
charged “bullets” passed
right through the gold
atoms in the sheet of
gold foil without changing
course at all.
⚫ Some of the positively
charged “bullets,”
however, did bounce
away from the gold sheet
as if they had hit
something solid. He
knew that positive
charges repel positive
charges.
⚫ This could only mean that the gold atoms in the
sheet were mostly open space. Atoms were not
a pudding filled with a positively charged
material.
⚫ Rutherford concluded that an atom had a small,
dense, positively charged center that repelled
his positively charged “bullets.”
⚫ He called the center of the atom the “nucleus”
⚫ The nucleus is tiny compared to the atom as a
whole.
Rutherford
⚫ Rutherford reasoned
that all of an atom’s
positively charged
particles were
contained in the
nucleus. The
negatively charged
particles were
scattered outside the
nucleus around the
atom’s edge.
NUCLEUS: The Center of an Atom
⚫ The NUCLEUS, that dense
central core of the atom,
contains both protons and
neutrons.
⚫ ELECTRONS are outside
the nucleus in energy
levels.
⚫ PROTONS have a positive
charge, neutrons have no
charge, and electrons have
a negative charge.
 The INADEQUECY of
Rutherford’s Model
JAMES MAXWELL “When a charged body moves
in an orbit around another
oppositely charged body, the
assembly of particles will
gradually lose energy by
emitting radiation. The loss of
energy would cause the
electrons to move inward and
circle even close to the
nucleus until eventually these
electrons would be indirect
contact with the nucleus.”
 IF MAXWELL
THEORY APPLIES
TO ATOMS THE
RUTHERFORD’S
ATOMS CANNOT BE
STABLE!!
 Bohr Model
⚫ In1913, the Danish
scientist Niels Bohr
proposed an
improvement. In his
model, he placed
each electron in a
specific energy
level.
Bohr Model
⚫ According to Bohr’s
atomic model,
electrons move in
definite orbits around
the nucleus, much
like planets circle the
sun. These orbits, or
energy levels, are
located at certain
distances from the
nucleus.
 Niels Bohr’s Model (1913)

⚫ Electrons orbit
the nucleus in
circular paths of
fixed energy
(energy levels).
 Bohr’s Postulates are as
follows:
1. A hydrogen atom consists of a nucleus
containing a proton and an electron. The
electron revolves around the nucleus in a
circular orbit. There is a force balances
the centrifugal force on electron.
 Bohr’s Postulates are as
follows:
2. Only certain circular orbits are permitted.
The energy of the electron in a given orbit is
fixed. As long as the electron stays in that
orbit, it neither absorbs nor radiates energy.
The non-radiating state is called the
STATIONARY STATE.
 Bohr’s Postulates are as
follows:
3. The electron may move from one
stationary state to another. To do so, it
must absorb or emit a quantity of energy
exactly equal to the difference in energy
between the two states.
 Bohr’s Model was important because it
introduced the idea of quantized energy
levels for electron’s in atoms. However, it
had many limitations.

IT COULD ONLY EXPLAIN ATOMS OR

IONS WITH A SINGLE ELECTRON SUCH
AS H, He, and Li. IT COULD NOT
EXPLAIN THE ATOMIC SPECTRA OF
ATOMS WITH MANY ELECTRONS.
 ATOMIC
SPECTRA
 Bohr's Model of the Atom

Niels Bohr (1913):

-studied the light produced when atoms
were excited by heat or electricity
 FLAME TEST
✓is a form of qualitative analysis that is
used to visually determine the identity
of an unknown metal or metalloid ion
based on the color
emission.
✓A distinctive color is emitted because
the heat of the flame excites the
electrons of the metal ions, causing
them to emit visible light.
 LIMITATIONS
➢ It cannot detect low concentrations of almost all
ions.
➢The intensity of the visible light differs from one
sample to another.
➢Contaminants affect the test results. Sodium, in
particular, is present in most compounds and will
color the flame. Sometimes a blue glass is used to
filter out the yellow of sodium.
➢The test cannot differentiate between all
elements. Several metals produce the
same flame color. Some compounds do not
change the color of the flame at all.
 ATOMIC SPECTRA
/ LINE SPECTRA
⚫ The spectra produced by some gaseous
substances consist of only a limited
number of colored lines with dark spaces
between them.
⚫ Each element has its own distinct line
spectrum, a kind of atomic fingerprint. It
can be used to identify an element.
SPECTROSCOPY
✓ the study of spectral
lines
✓ was developed by
Robert Bunsen and
Gustav Kirchoff
 SPECTROSCOPE
⚫ Its an instrument that is used for analyzing
colors given off by the vapors of elements

A glass prism separatess the light given off into its

component wavelength. The spectrum produce appears as a
series of sharp lines wih characteristics of colors and
wavelenght on a dark background instead being continous
like the rainbow.. The color, number and position of lines
produced is called the “FINGERPRINT” of an element.
 Niels Bohr’s Atom
⚫ Electronscan jump from energy level
to energy level.

⚫ Electronsabsorb or emit light energy

when they jump from one energy level
to another.
GROUND STATE – when electrons normally exists
in the lowest energy state.
EXCITED STATE – when an electron goes into
higher energy state.

EXCITATION IS ACHIEVED BY
SUPPLYING ENERGY TO THE ATOM
 Quantum
⚫A quantum of energy is the amount
of energy required to move an
electron from one energy level to
another.
The energy levels are like the rungs
of a ladder but are not equally
spaced.
 Photons

⚫ Photons are bundles of light energy that

is emitted by electrons as they go from
higher energy levels to lower levels.
Bohr Model of Atom Increasing energy
n=3
of orbits

e-
e- n=2

e-
n=1
e-
e-
e- e- e-

e-
e-
e-

A photon is emitted
with energy E = hf
 IN THE
ORIGINAL
BOHR’S
MODEL
 Niels Bohr’s Model (1913)

⚫ Electrons orbit
the nucleus in
circular paths of
fixed energy
(energy levels).
A circular orbit
was specified by
a whole number

n with values
1,2,3 and so
on. This is
referred as the
PRINCIPAL
QUANTUM
NUMBER.
ARNOLD SOMMERFIELD
✓Introduced the concept of elliptical orbits
to explain the splitting of spectral lines.
✓This becomes the second quantum of
energy sublevels.

s – SHARP
p- PRINCIPAL
d- DIFFUSE
f- FUNDAMENTAL LINES
ADDITIONAL FEATURES OF BOHR’S
MODEL
✓ As the principal quantum number (n) INCREASES, the
orbits (shells) become LONGER (radius increases).
✓Electrons in n=1 orbit/ shell have the lowest energy. As the
value of n increases, the energy of the electron increases.
✓Each orbit/ shell can hold a maximum of 2n2 electrons.

orbit maximum # electrons

1 2
2 8
3 18
4 32
 Bohr's Model of the Atom

e.g. fluorine:
#P =

#e- =

#N =
 Bohr's Model of the Atom

e.g. fluorine:
#P = atomic #
=9
#e- =

#N =
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = # P
=9
#N =
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9

#N = atomic mass - # P
= 10
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10

draw the nucleus with

protons & neutrons

 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10
how
many electrons can
fit in
the first orbit?
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10
how many electrons can
fit in the first orbit?
2
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10

how many electrons are left?

 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10

how many electrons are left? 7

 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10

how many electrons are left? 7

how many electrons fit in the
second orbit?
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10

how many electrons are left? 7

how many electrons fit in the
second orbit? 8
 Bohr's Model of the Atom

e.g. fluorine:
#P = 9

#e- = 9 9P
10N
#N = 10
REMEMBER!
The Bohr model explained
the emission spectrum of
the hydrogen atom but did
not always explain those
of other elements.
 Quantum Mechanical Model

⚫ 1920’s
⚫ Werner Heisenberg (Uncertainty Principle)
⚫ Louis de Broglie (electron has wave
properties)
⚫ Erwin Schrodinger (mathematical equations
using probability, quantum numbers)
Werner Heisenberg: Uncertainty Principle

⚫ We can not
know both the
position and
momentum of
a particle at a
given time.
Louis de Broglie, (France, 1892-1987)
Wave Properties of Matter (1923)
⚫Since light waves have a
particle behavior (as
shown by Einstein in the
Photoelectric Effect), then
particles could have a
wave behavior.
⚫de Broglie wavelength
l= h
mv
Electron Motion Around Atom
Shown as a de Broglie Wave
Davisson and Germer (USA, 1927)
confirmed de Broglie’s hypothesis
for electrons.

Electrons produced a diffraction

pattern similar to x-rays.
 Erwin Schrodinger, 1925
Quantum (wave) Mechanical Model
of the Atom
⚫ Four quantum
numbers are
required to
describe the state
of the hydrogen
atom.
Atomic Orbital s

2s
 The 3 p orbitals

http://www.rmutphysics.com/CHARUD/scibook/crystal-structure/porbital.gif
⚫ The d orbitals
f orbitals
 ELECTRON
CONFIGURATION