Download History of the Atom notes

Unit 3
History of the Atom
Democritus
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about 460 bc
stated that the “ultimate particle” is
atmos
He believed that atoms were indivisible
and indestructible
was not based on the scientific method
– but just philosophy
John Dalton - 1803
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English school teacher
the atom was a small, indivisible
particle like a “small marble”
Dalton’s Atomic Theory (experiment based!)
John Dalton
(1766 – 1844)
1) All matter is made of tiny
indivisible particles called
atoms
2) Atoms of the same element are
chemically the same. Atoms of
different elements are
chemically different.
3) Individual atoms of an element have slightly
different masses. We use average mass.
4) Different elements have different ave masses
5) Atoms can’t be divided in normal chemical
reactions.
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray
tube to deduce the presence of a negatively
charged particle: the electron
Modern Cathode Ray Tubes
Television
Computer Monitor
Cathode ray tubes pass electricity
through a gas that is contained at a
very low pressure.
Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons
were like plums embedded in a
positively charged “pudding,” thus it
was called the “plum pudding” model.
Ernest Rutherford’s
Gold Foil Experiment - 1911
Alpha particles are helium nuclei The alpha particles were fired at a thin
sheet of gold foil
 Particles that hit on the detecting
screen (film) are recorded
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Rutherford’s Findings
Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
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“Like howitzer shells bouncing
off of tissue paper!”
Conclusions:
a) The nucleus is small
b) The nucleus is dense
c) The nucleus is positively
charged
The Rutherford Atomic Model
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Based on his experimental evidence:
 atom - mostly empty space
 all positive charge, and most mass is in
the center called a “nucleus”
 nucleus = protons and neutrons
(neutrons confirmed by Chadwick in
1932)
 electrons distributed around the
nucleus, and occupy most of the volume
 His model was called a “nuclear model”
Niels Bohr - 1913
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Niels Bohr wondered why horseshoes
heated in a forge changed color, but did
not react
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Nucleus?
Electrons?
http://www.askaboutireland.ie/_internal/gxml!0/m6s6aicjogoampxnh1a9dzzq7k4g1nn$o0srg946q
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Neils Bohr
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All atoms contain energy – the energy
has to do with the electrons
1913 Bohr said that electrons are in
definite areas with definite amounts of
energy
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Planetary Model
Electrons are in the
“ground state” – normal
lowest energy state
*excited state electrons
gain energy in fixed
amounts (photons) and
go to higher levels
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E = hf
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E = energy
h = Planck’s constant
f = frequency
So, Energy = Planck’s constant x frequency
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Electrons don’t stay
in the excited state,
but fall back to the
ground state and
give off energy in
the form of light
called a spectrum
(which is unique for
each element)
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http://www.mhhe.com/physsci/astrono
my/applets/Bohr/applet_files/Bohr.html
http://www.colorado.edu/physics/2000/
quantumzone/bohr.html
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Speed of light (constant) = 2.998 x 108 m/sec
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c = fλ
speed
of light
frequency
wavelength,
lambda
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c = fλ
frequency=speed/wavelength
2.998 x 108 m/s divided by 7.6 x 10-7
= 3.9 x 1014 1/s = 3.9 x 1014 Hz
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Energy of an electron = 2.179 x 10-18 J/n2
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Schrodinger - 1925
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Wave Mechanical Model – electrons
exist in areas of probability called
“space orbitals” (Rutherford and Bohr
support this)
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Quantum Mechanics – each electron is
identified by 4 quantum numbers:
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1.
2.
3.
4.
Principle quantum # - level
Orbital quantum # - shape
Magnetic quantum # - orientation
Spin quantum # - clockwise or counter
http://www.colorado.edu/physics/2000/
quantumzone/schroedinger.html
Werner Karl Heisenberg -1927
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Uncertainty Principle – you can’t locate
the exact position of an electron at any
given time (too small, too fast)