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ATOMIC THEORY
Early History of
Atomic Theories
Greek Philosophers
Empedocles

Proposed that all matter was made
up of four basic particles:
 Air
 Earth
 Fire
 Water
Democritus

Proposed that matter was
made up of tiny indivisible
particles
 These particles could not be
seen by the naked eye and
could not be cut or
subdivided into smaller bits
 They were called atoms
Dalton Atomic Theory

John Dalton in 1805 recreated the modern theory of
atoms
 Explained three important
scientific laws
Dalton’s Laws
 Law
of definite composition: each
element has a particular combining
capacity
 Law of multiple proportions: some
atoms have more than one combining
capacity
 Law of conservation of mass: atoms are
neither altered, created nor destroyed in
a chemical reaction
Dalton Atomic Theory

Matter is composed of
indestructible, indivisible
atoms, which are identical
for one element, but
different from other
elements.
 Model: An atom is a solid
sphere like a marble of
billiard ball
J.J. Thomson’s Atomic Theory
1880's

Experiments with cathode rays resulted in
the discovery of the electron, a negatively
charged particle
 Cathode Ray Animation
Thomson Atomic Theory 1897

Matter is composed
of atoms that
contain electrons
(negative particles)
embedded in a
positive material
 Model: Raisin bun or
Plum Pudding
Rutherford Atomic Theory

Gold Foil Experiment
Rutherford Atomic Theory

An atom consists of a
small nucleus, which
contains positive charges
 Negative electrons circle
the nucleus in orbits
 Most of the atom consists
of empty space
 Model: Beehive
Rutherford Atomic Theory

An atom is made up of an equal number of
negatively charged electrons and positively
charged protons
 Most of the mass of the atom and all of its
positive charge is contained in a tiny core
region called the nucleus
 The nucleus contains protons and neutrons
(Chadwick, 1932) that have approximately
the same mass
 The number of protons is the atomic number
(Z)
 The total number of protons and neutrons is
called the mass number (A)
Isotopes

It has been discovered that atoms of an
element can have different numbers of
neutrons.
 These different varieties of an element are
known as isotopes.
 For example, hydrogen has three naturallyoccurring isotopes: a hydrogen atom with
one proton and one electron, a second
isotope with one proton, one neutron and one
electron and the third isotope has one proton,
two neutrons and one electron.
 We signify these different isotopes using
isotope notation.
Isotope Notation
1
1
H
2
1
H
3
1
H

The three numbers in the lower lefthand corner are the atomic
numbers of the hydrogen atoms
 In each case, the atomic number is
1
 In the top left corner is the mass
number, which is defined as the
sum of the protons and neutrons in
the nucleus
A General Form for Isotope
Notation
A
Z
X
Where Z is the atomic number
and A is the mass number
Atomic Spectra
 Flames
and other glowing gases
produce discontinuous spectra called
line spectra, which consist of distinct
lines
 Excited elements emit only certain
wavelengths of radiation, which are
unique for each element
 This distinct pattern can be used to
identify individual atoms. This practice
is termed SPECTROSCOPY
Spectroscopy
 Spectroscopy
is an analytic method that
uses a spectroscope to separate the
entering light into its components by
means of a prism or a diffraction grating
Bohr’s Model
 Electrons
are restricted
to specific energy levels
 Each level represents
an increase in energy
 Electrons can jump
energy levels if
“excited” called an
electron transition
 Energy is released as
the electrons fall back
to their ground state
 This
is an example
of discontinuous
spectra or line
spectra
 Line spectra is
unique to each
element