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The History of
Atomic Theory
A long and winding road
Section 4.1
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The “Atom”
 The smallest particle of an element
that retains its identity in a chemical
reaction.
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Aristotle
 Matter is in continuous motion
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Democritus
 Greek Philosopher (460 BC- 370 BC); the
“laughing philosopher”
 Atoms are indivisible and indestructible
 Did not explain chemical behavior and
lacked experimental support.
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Dalton’s Atomic Theory
John Dalton (1766-1844)
1.
All elements are composed of tiny
indivisible particles called atoms.
2.
Atoms of the same element are identical.
The atoms of any one element are different
from the atoms of any other element.
3.
Atoms of different elements can physically
mix together or can chemically combine in
whole-number ratios to form compounds.
4.
Chemical reactions occur when atoms are
separated, joined, or rearranged. Atoms of
one element, however, are never changed
into atoms of another element as a result of
a chemical reaction.
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Joseph John Thomson (1856-1940)
 Adds detail to Dalton’s theory.
 “Plum pudding model”- negatively charged subatomic
particles within positively charged atom.
 Called particles “corpuscles”
 Cathode ray tube
 electrons
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Robert A. Millikan
Nobel Prize in 1923
 Conducted an experiment to find the quantity of charge
carried by an electron
 Oil-drop experiment
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Rutherford Model (1911)
“the father of nuclear physics”
 Gold Foil Experiment- positively
charged alpha particles aimed at
thin sheet of gold foil.
 Proposed that atom is mostly
empty space, with positive charge
concentrated in the center.
Electrons move around the nucleus.
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To sum up Rutherford’s
Gold Foil Experiment
RESULTS
CONCLUSIONS
 Not a lot of charged
particles located
particles pass
throughout atom
through the gold foil
 Particles ran into
 Few alpha particles
small amount of
were deflected (went
charge
off at an angle)
 Strong concentration
 Very few alpha
of positively charged
particles were
particles (nucleus)
reflected right back
 Most of the alpha
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James Chadwick (1891-1974)
 February 1932- “The possible existence of a neutron”
 May 1932- “The existence of a neutron”
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The Bohr Model
Neils Bohr- 1913
 Proposed that an electron is
found in specific circular
paths, or orbits, around the
nucleus.
 Each orbit has a fixed energy.
 A quantum of energy is the
amount of energy required to
move an electron from one
energy level to another.
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Quantum Mechanical Model
(WAVE-MECHANICAL)
 1926- Erwin Schrodinger: mathematical equation
describing the behavior of the electron in a hydrogen
atom.
 This model determines the allowed energies an electron
can have and how likely it is to find the electron in various
locations around the nucleus.
 “electron cloud”
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Atomic Orbitals
 A region of space in which there is a high probability of
finding an electron.
 Each energy sublevel corresponds to an orbital of different
shape describing where the electron is likely to be found.
 More on this later….
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Atomic Models…a recap
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Atomic Details
Sections 4.2 and 4.3
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Subatomic Particles
Particle
Symbol
Relative
Charge
Relative
Mass
Actual
Mass (g)
Electron
e-
1-
1/1840
9.11 x 10-28
Proton
p+
1+
1
1.67 x 10-24
Neutron
n°
0
1
1.67 x 10-24
Atomic Number
 Elements are different because they contain a different
number of protons
 Atomic # = # of p+ in the nucleus of an atom of that
element
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Atoms are
 ELECTRICALLY NEUTRAL
 # of protons = # of electrons

In an electrically neutral atom:
 Atomic # = # of electrons = # of protons
 Protons = Electrons = Atomic number
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Mass Number
 Most of the mass of an atom is concentrated in the
nucleus
 Mass Number = # of p+ + # of no
 Number of Neutrons= Mass # - Atomic #
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What does the Periodic Table tell us
about subatomic particles?
Mass Number
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C
Element Symbolone or two letter
abbreviation of element
name
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Atomic Number = # of
protons
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Isotopes
 Atoms that have the same number of protons but
different numbers of neutrons.
 Because isotopes of an element have different
numbers of neutrons, they also have different mass
numbers
 Ex. C-12 and C-14
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Atomic Mass
 Most elements occur as a mixture of two or
more isotopes, each with a fixed mass and a
natural % abundance.
 ATOMIC MASS= the weighted average mass of
the atoms in a naturally occurring sample of the
element.
 To calculate the atomic mass of an element,
multiply the mass of each isotope by its natural
abundance (expressed as a decimal) and then add
the products.
Calculate the average atomic
mass of Sc if:
 Sc-44: 50% abundance
 Sc-45: 15% abundance
 Sc-46: 35% abundance
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Calculate the Average Atomic
Mass of copper:
69.1% Cu-63
30.9% Cu-65
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Calculate the Average Atomic
Mass of boron:
19.9% B-10
80.1% B-11
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MOLE
 SI Unit
 Abbreviated mol
 Amount of a substance that contains the same number of
particles as the number of atoms in exactly 12 g of C-12
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6.02 x 10
 Avogadro’s Number
 Number of particles in exactly one mole of a pure
substance
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Molar Mass
 Gram Atomic Mass (g.a.m.)
 Mass (g) of one mole of a pure substance
 Mass of 1 mole atoms (g) = atomic mass (amu)
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Conversion Factors
 1 mol = 6.02 x 1023
 1 mol = mass (g) from P. Table
 Most atomic masses are known to 4 or more sig figs
 For calculations, we will round the atomic masses to the tenths
place
• 1 mole C = 6.02 x 1023 atoms of C
• 1 mole C = 12.0 g
• 1 mole Cu = 6.02 x 10 23 atoms of Cu
• 1 mole Cu = 63.5 g
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Conversion Problems
 Grams  MOLE  Atoms
 What is the mass of 2.8 mol He?
 How many moles are represented by 7.11 x 1024 atoms of
Hg?
 A sample containing 59.2 g of Al contains how many
atoms?
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