Ch. 3: Development of Atomic Theory
Democritus
 John Dalton
 J.J. Thomson
 Robert A Millikan
 Lord Kelvin
 Rutherford, Geiger
& Marsden
or “What does an atom look like?”

Democritus, 400 BC
“Everything is made up of a few
simple parts called atomos.”
Atomos means “uncuttable” in
Greek.
 He envisioned atomos as small,
solid particles of many different
sizes and shapes.
 *His ideas were rejected later
by Aristotle who supported the
“earth, air, water, and fire”
concept of matter.

Laws of 1790’s

Law of conservation of mass: mass is neither
created nor destroyed during ordinary
chemical or physical reactions.
Law of Conservation of Mass


Law of Conservation of Mass
In a combustion reaction, 46.0 g of ethanol reacts with 96.0 g
of oxygen to produce water and carbon dioxide. If 54.0 g of
water is produced, then how much carbon dioxide is produced?
Law of Definite Proportions or
Constant Composition

Law of definite proportions: A pure
compound always contains the same
elements in the same proportions by
mass.

Ex. Law of Constant Composition


A sample of chloroform is found to contain 12.0g of carbon,
106.4g of chlorine and 1.01g of hydrogen. If a second
sample of chloroform is found to contain 30.0 g of carbon,
how many g of chlorine and hydrogen does it contain?
Law of Multiple Proportions
Ex. Law of multiple proportions:When
elements combine to form more than a
single compound, the ratios of the masses of
the combining elements can be expressed by
a ratio of small whole numbers.
Law of Multiple Proportions
Law of Multiple Proportions
Water contains 2.02g of hydrogen and 16.0g of oxygen.
Hydrogen peroxide( H2O2) contains 2.02 g of
hydrogen and 32.0g of oxygen. Show how these data
illustrate the law of multiple proportions?
Animation:
http://cwx.prenhall.com/petrucci/medialib/media_portfolio/02.html
http://chemsite.lsrhs.net/c_AtomicTheory/dalton.html
http://books.nap.edu/books/030907309X/html/images/p2000472eg16001.jpg
John Dalton, 1808
http://www.imageil.com/html_ver/cbt/dalton/dalton.swf
The main ideas of his theory
are:
•Elements are made of tiny
particles called atoms.
All atoms of an element are
identical.
•Atoms of a given element are
different from those of the
other elements.
•Atoms of one element can
combine with atoms of other
elements to form compounds. A
given compound always has the
same relative number and types
of atoms.
Atoms are indivisible in the
chemical process. A chemical
reaction simply changes the way
the atoms are grouped together.
 Dalton’s theory successfully
explained ‘law of conservation of
mass’ ‘ law of definite
proportions’ and the ‘law of
multiple proportions’. How?

J. J. Thomson (Joseph John Thmompson),1897
http://physics.nad.ru/Physics/Oscil2.gif

Discovery of electrons (J. J.
Thomson’s cathode ray tube
experiment): J. J.Thomson
showed in late 1890s that
atoms of any element can be
made to emit tiny negative
particles that he called as
“electrons”. He conducted
his experiments using
cathode ray tube (CRT). He
concluded that all types of
atoms must contain this
particle, which he called as
‘electron’.
On the basis of his experiments he further predicted that the
atoms should also have a positively charged particle to balance
out the negative charge of the electrons, since the atom is
electrically neutral. He also calculated charge to mass ratio
(e/m) for electrons.
http://chemsite.lsrhs.net/c_AtomicTheory/Images/overheads_screen/JJThompson_lage.jpg
http://www.lip.pt/~outreach/experiments/f+0+0.jpg
Robert A. Millikan, 1909
Calculated the charge on the electron.
 Electrons are present in atoms of all
elements.
It was inferred that:
 atoms contain a positive charge to balance
the negative electrons.
 Atoms contain other particles that account
for most of the mass.

http://www.bun.falkenberg.se/gymnasium/amnen/fysik/millikaneng.html
http://physics.nad.ru/Physics/mill.gif
http://chem100a-9.chem.lsu.edu/matter/chap26/animate1/an26_003.mov
.
Milliken’s oil drop experiment: American scientist
Robert Milliken calculated the charge of electron (e) by
conducting his famous oil drop experiment.
http://micro.magnet.fsu.edu/optics/timeline/people/antiqueimages/millikan.jpg
Millikan Oil Drop Experiment
Robert Millikan
(University of Chicago)
determined the charge
on the electron in 1909.
ANIMATION:
http://physics-animations.com/Physics/English/top_ref.htm
© 2009, Prentice-Hall,
Inc.
Lord Kelvin’s Plum Pudding Model
Plum Pudding Model : Lord
Kelvin proposed a plum
pudding model for the
structure of atom in 1910. In
this model the negatively
charged electrons were
pictured as embedded in a
positively charged spherical
cloud much as raisins are
distributed in an old fashioned
plum pudding.
Lord Kelvin (William Thompson),1910


Proposed a plum pudding
model
The negatively charged
electrons were pictured as
embedded in a positively
charged spherical cloud
much as raisins are
distributed in a plum
pudding.
Goldstein’s Canal Ray Experiment
Goldstein discovered protons by using
the same apparatus that was used for
cathode ray discovery.
 He used a perforated anode in the CRT
and found that there were these
positively charged rays (canal rays), which
consisted on positively charged particles
called as protons.

Discovery of Protons
In 1886, Goldstein discovered positively charged particles in
atom, which he called as “protons”.
http://images.google.com/images?hl=en&q=Goldstein's%20and%20discovery%20of%20protons&um=1&ie=UTF-8&sa=N&tab=wi
Rutherford’s Gold Foil Experiment
In 1911, Ernset Rutherford carried out his famous gold
foil experiment changing the idea about the structure of
atom dramatically. This experiment included bombarding a
very thin layer of gold foil with alpha particles that are
positively charged. Surrounding the foil was alpha particle
detector that glowed each time an alpha particle hit it.
Although most of the alpha particle passed straight through
the foil, very few particles deflected and a very small number
bounced back. Based on his experiment Rutherford
concluded that
•Most of the space in the atom is empty.
•In the center of the atom is a very small and positive
nucleus.
Thus, Rutherford came up with the idea of a ‘nuclear atom’.
http://images.google.com/imgres?imgurl=http://www.newgenevacenter.org/portrait/rutherford.jpg&imgrefurl=http://www.newgenevacenter.org/sci3_quantum2.htm&h=218&w=154&sz=6&tbnid=98HQ
ryBKGdkd_M:&tbnh=102&tbnw=72&prev=/images%3Fq%3Drutherford&start=3&sa=X&oi=images&ct=image&cd=3
Rutherford’s Gold Foil Experiment Contd.
http://www.chemsoc.org/timeline/pages/timeline.html)
Rutherford, Geiger & Marsden,1911



Discoved that the mass of
the atom is contained in a
tiny positive nucleus
Inferred most of the atom
is empty space.
Discovered nucleus by
shooting positively charged
alpha particles at a thin
piece of gold foil and
observing how some, but
very few, of the particles
were deflected.
http://www.mhhe.com/physsci/chemis
try/essentialchemistry/flash/ruther14.
swf
Animation:
http://cwx.prenhall.com/petrucci/medialib/media_portfolio/02.html
Discovery of Neutrons (James
Chadwick) 1932
http://www.google.com/search?hl=en&lr=&q=picture+of+James+chadwick
In 1932, James Chadwick discovered that most
nuclei also contain another neutral particle
called neutron, which is slightly more massive
than proton but has no charge. His experiment
involved bombarding Beryllium atoms with alpha
particles, which produced a strong beam of particles
that was not deflected by electrical field.
James Chadwick, 1932


In the nucleus, there is also a particle
without charge, called a neutron,
which is about the size of a proton
and has a slightly greater mass.
Chadwick bombarded Beryllium
atoms with alpha particles which
produced a strong beam of particles
that were not deflected by an
electrical field.
Bohr’s Model : Planetary Model of Atom
http://web.gc.cuny.edu/ashp/nml/copenhagen/Bohr.jpg
Niels Bohr’s theory
stated that the electrons
move around the nucleus
in circular orbits with
certain energy, just as
the planets move around
the sun. Bohr’s most
important contribution
was the concept that the
energy of these orbits is
quantized.
Niels Bohr, 1913
His
model of the hydrogen atom consists of electrons
circling the nucleus but are restricted to particular orbits,
like the planets around the Sun.
The
electron is in its lowest energy state when it is in
the orbit closest to the nucleus.The energy is higher
when its in orbits successively farther from the nucleus.
It
can move to a higher energy orbit by gaining an
amount of energy equal to the difference between the
higher energy orbit and the initial energy orbit.
When
an electron drops back to a lower energy orbit a
photon is emitted that has energy equal to the energy
difference between the the two orbits.
Bohr’s Model
The orbits next to the nucleus have the least energy
and the one farthest from the nucleus has the most
energy. When an electron jumps from lower energy orbit
to a higher energy one, it absorbs energy and when it
moves form a higher to lower energy orbit it gives out
energy. This energy exchange also is quantized.
The shortcoming of Bohr’s model was that it could
only explain the atomic spectrum of hydrogen, but
could not explain the atomic spectrum of other
elements.
The Nature of Energy
Electrons in an atom can only
occupy certain orbits
(corresponding to certain
energies).
© 2009, Prentice-Hall, Inc.
Erwin Schrodinger, 1926




Created a mathematical model that
described electrons as waves.
Proved that the electrons do not orbit the
nucleus. Instead, their position is
determined by probability. 95% of the time,
they can be found in Bohr’s proposed orbits.
They exist in orbitals: a 3 dimensional
region around the nucleus that indicates the
probable location of an electron.
http://www.msu.edu/~russe153/schrodinger(young).jpg
Modern Concept of Atomic Structure(Wave
Mechanical Model) by Schrodinger:
Atom is divisible and has three subatomic particleselectrons, protons and neutrons, which further are
made up of sub sub atomic particle (such as mesons,
antineutrino etc.). Protons and neutrons are present in the
nucleus while the electrons move around the nucleus in an
electron cloud. Electron cloud is the place around the
atom where the probability of finding the electron is
the most. (>90%).
 Atom has mostly empty space, with a small positive
nucleus in the center. (If nucleus is the size of the
grape, the electrons would be about a mile away
from the nucleus.)
 If all the atoms are made of three basic subatomic particles,
then why are atoms different?
http://www.msu.edu/~russe153/schrodinger(young).jpg
http://www.epa.gov/radiation/graphics/shrodinger.jpg

http://www.msu.edu/~russe153
/schrodinger(young).jpg
Five Major Atomic Models
To sum it all, there are five major
atomic models:
1. Dalton’s Atomic Model 1808
2. Plum Pudding Model
1909
3.Rutherford’s Nuclear Model 1910
4. Bohr’s Solar System Model 1913
5. Shrodinger’s Wave Mechanical Model
1927

Subatomic Particles (SAP’s)
Particle & Symbol
Proton
(p+, 11H)
# determines elements identity
Neutron (n0,01n)
Electron
(e-, -10e)
Relative
Electrical
Charge
+1
positive
0
uncharge
d
-1
negative
Relative
Mass
Actual Mass (kg)
1.673*10-27
1.007
1836 times greater than
the mass of an e-
1.009
1.675*10-27
0.0005
9.109*10-31
Atomic Vocabulary
Mass number = p + n (in the nucleus)
Element
symbol
A
Z
E
Atomic number = # p (=#e-)
◦ Elements are put in order of increasing atomic
number on the periodic table, identifies an element.
Ex: An atom of carbon with 7
neutrons:
13C
6
207Pb
82
Isotopes
Isotopes: atoms of one element that have different mass
numbers (different # of n)
Ex: There are 3 naturally occurring isotopes of hydrogen:
Isotope
Name
#p
Protium
Deuterium
Tritium
#n
% of all H
1
0
99.985
1
1
0.015
1
2
< 0.001
Relative Atomic Mass
Relative Atomic Mass: It is convenient
to use relative atomic mass because the
atoms are too small. (Oxygen atom
weighs 2.657X 10-23g).
 The atomic mass of any nucleide is
determined by comparing it with carbon12 atom.
 One atomic mass unit (amu) is exactly
1/12 of the mass of one C-12 atom.
(1.660 540 X 10-27 kg)

Average Atomic Mass
Average atomic mass: It is the
weighted average of the atomic masses of
naturally occurring isotopes of an
element.
 Formula for calculating average
atomic mass
 (mass of nucleide-1 X decimal
fraction for its % in the mixture) +
(mass of nucleide-2 X decimal
fraction for its % in the mixture)

Ex for Calculating Avg. Atomic Mass
Carbon exists as two isotopes, C-12 and C-13.
Their percent abundances are 89.892% and
1.108%. Calculate Carbon’s average atomic
mass.
Average atomic mass (or atomic weight):
12C: 98.892% x 12 amu
=11.867
+
13C: 1.108% x 13.00335 amu = 0.1441
AW = 12.011 amu…....
Therefore 12.011 g C = 1 mol C
The MOLE
The mole: One mole of something
contains 6.022 X1023 units of that
substance.
 Ex. One mole of Hydrogen atoms
contains exactly 6.022 X1023 atoms
of Hydrogen.
 Ex. H2O has 1 mole of water
molecules, 1 mole of oxygen atoms,
2 moles of hydrogen atoms and 3
moles of atoms in all.

Avogadro’s Number
Avogadro’s number: It is defined as
the number of particles in one mole
of any pure substance.This number
is found to be 6.022 X1023.
Molar Mass
Molar mass: Mass of one mole of a
substance. Molar mass can be
calculated by expressing the mass in
amu in g.
 Molar mass of C = 12 g

Avogadro’s Number
6.02 x 1023
 1 mole of 12C has a
mass of 12 g.

© 2009, Prentice-Hall,
Inc.
Molar Mass

By definition, a molar mass is the mass of
1 mol of a substance (i.e., g/mol).
◦ The molar mass of an element is the mass
number for the element that we find on the
periodic table.
◦ The formula weight (in amu’s) will be the
same number as the molar mass (in g/mol).
©
200
9,
Pre
ntic
eHall,
Inc.
Using Moles
Moles provide a bridge from the molecular scale
to the real-world scale.
© 2009, Prentice-Hall,
Inc.
Mole Relationships
One mole of atoms, ions, or molecules contains
Avogadro’s number of those particles.
 One mole of molecules or formula units contains
Avogadro’s number times the number of atoms or
ions of each element in the compound.

© 2009, Prentice-Hall,
Inc.