the development of atomic theory
TRANSCRIPT
The Development of Atomic Theory
SCH12U
February 3 2011
Mr. Dvorsky
• Nearly 2500 years ago Greek philosophers (i.e. Democritus) expressed a belief matter is composed of tiny indivisible particles called atoms (atomos is the Greek word for “indivisible”)
• These conclusions were not based on any evidence; they were derived from philosophical reasoning.
• Experimentation by many scientists during the 18th and 19th centuries led to the development of 2 Laws:
1. The Law of Conservation of Mass
• During chemical change no loss or gain of mass occurs.
2. The Law of Definite Proportions
• Compounds contain elements in fixed proportions by mass.
John Dalton, early 19th century took these findings
and developed Atomic Theory.
1. Matter consists of particles called atoms.
2. Atoms are indestructible. In chemical reactions atoms rearrange but are not broken apart.
3. Atoms in one particular element are identical, but differ from atoms of other elements.
4. Compounds are created when atoms of different elements combine in definite proportions.
Source ofElectricalPotential
Metal Plate
Gas-filledglass tube Metal plate
Stream of negativeparticles (electrons)
JJ Thompson –
Discovery of Electrons
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58
A Cathode Ray Tube
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58
Thomson’s Experiment
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vacuum tube
metal disks
voltage source
Thomson’s Experiment
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Passing an electric current makes a beam appear
to move from the negative to the positive end
Thomson’s Experiment
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Thomson’s Experiment
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By adding an electric field…
he found that the moving pieces were negative.
The Effect of an Obstruction on
Cathode Rays
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 117
The Effect of an Electric Field on
Cathode Rays
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 118
J.J. Thomson
• He proved that atoms of any element can be made to emit tiny negative particles.
• From this he concluded that ALL atoms must contain these negative particles.
• He knew that atoms did not have a net negative charge and so there must be balancing the negative charge.
J.J. Thomson
William Thomson (Lord Kelvin)
• In 1910 proposed
the Plum Pudding
model
– Negative electrons
were embedded into
a positively charged
spherical cloud.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
Spherical cloud ofPositive charge
Electrons
Plum-Pudding Model
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
Thomson Model of the Atom
• J.J. Thomson discovered the electron and knew that electrons could be emitted from matter (1897).
• William Thomson proposed that atoms consist of small, negative electrons embedded in a massive, positive sphere.
• The electrons were like currants in a plum pudding.
• This is called the ‘plum pudding’ model of the atom.
electrons-
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Rutherford, see animation
-fired alpha particles at a very thin piece of foil.
-The alpha particles to pass through without
changing direction (very much)
Because....
-The positive charges were spread out evenly.
Alone they were not enough to stop the alpha
particles
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Because he thought the mass was evenly distributed in the atom.
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Because, he thought the mass was evenly distributed in the atom
Explanation of Alpha-Scattering Results
Plum-pudding atom
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Alpha particles
Nuclear atom
Nucleus
Thomson’s model Rutherford’s model
The Rutherford Atom
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 323
• Bohr suggested the planetary model of the atom could be rescued if one assumption is made: certain special “states of motion” of the electron corresponding to electron shells would not result in radiation and therefore the electron can exist indefinitely.
Bohr was saying, in effect, is that the atom can exist
only in certain discrete energy states: the energy of
the atom is quantized. Bohr noted that this
quantization nicely explained the observed emission
spectrum of the hydrogen atom. The electron is
normally in its smallest allowed orbit, corresponding
to n = 1; upon excitation in an electrical discharge or
by ultraviolet light, the atom absorbs energy and
the electron gets promoted to higher quantum levels.
These higher excited states of the atom are
unstable, so after a very short time (around 10—
9 sec) the electron falls into lower orbits and
finally into the innermost one, which
corresponds to the atom's ground state. The
energy lost on each jump is given off as a
photon, and the frequency of this light provides
a direct experimental measurement of the
difference in the energies of the two states,
according to the Planck-Einstein
relationship e = hν.
