Transcript Atomic models 300

Atomic Models
Atomic Theories
Atomic Theory – A Short History
Fifth Century, BCE
Democritus
Believed matter was composed of very
small, individual particles that were
indestructible
He called them “atomos” (meaning
uncuttable)
His ideas persisted for centuries even
though there was no experimental proof
Atomic Theories
JOHN DALTON - 1808
Revised early Greek ideas
into testable scientific theory
Based his Atomic Theory on
three important concepts:
1. Law of Conservation of Mass
2. Law of Multiple Proportions
3. Law of Definite Proportions
Atomic Theories
Law of Conservation of Mass
States that mass cannot be created or
destroyed
the mass of the reactants equals the mass of the
products
Atomic Theories
Dalton’s Principles
1. All matter is composed of extremely small
particles called atoms which cannot be
subdivided, created or destroyed
2. Atoms of a given element are identical in
their physical and chemical properties
Example: all water molecules freeze at 0 deg C
and react with explosively with sodium
Atomic Theories
Dalton’s Principles (continued)
3. All atoms of one element are different
from those of any other element
4. Atoms combine in simple, wholenumbered ratios to form compounds
Based on the Laws of Definite and Multiple
Proportions
Atomic Theories
Dalton’s Principles (continued)
5. In chemical reactions, atoms are
combined, separated or rearranged but
NEVER created, destroyed or changed
Based on The Law of Conservation of Mass
Atomic Theories
Dalton, however, did all this work in the
early 1800’s without ever knowing about
subatomic particles!
(Protons, Neutrons, Electrons)
Atomic Models
The Adventures of J.J. Thomson, Plum
Pudding and the Electron!
MMMMM…..That
plum pudding
looks delicious!
Atomic Models
JJ Thomson:
 English Physicist
 Experimented with “cathode
rays”
 Was able to determine that
the mass of the particles in
the ray was much smaller
than the mass of a
hydrogen atom
The particle must be
a smaller than an
atom!
I LOVE
plum
pudding!
Atomic Models
JJ Thomson’s Cathode Ray Tube (CRT)
Anode – attached to the
positive terminal of the
voltage source
Cathode – attached to the
Negative end of the voltage
source
Atomic Models
A magnet placed near the tube deflected the
beam, proving it was negatively charged.
A small paddlewheel in the tube turned when
hit by the beam, meaning the particles had
mass.
Atomic Models
Thompson called the negatively charged
particles in the beam “corpuscles”
This name was later changed to
“electrons” by one of Thompson’s
associates.
Atomic Models
Thompson’s Conclusion:
An atom is a sphere of positive
charge with electrons
embedded in it
“Plum-Pudding”
Model of the
Atom
Atomic Models
In 1909, one of his students, Ernest
Rutherford, disproved the “Plum
Pudding” model by doing is famous
“Gold Foil” experiment.
Atomic Models
Rutherford’s experiment
http://www.mhhe.com/physsci/chemistry/essen
tialchemistry/flash/ruther14.swf
Atomic Models
• What conclusions were drawn from
Rutherford’s Experiment?
Atomic Models
Rutherford’s Model of the Atom:
Electrons orbit the nucleus just as planets
orbit the sun
Did not explain why the negatively charged
electrons did not crash into the
positively charged nucleus.
Atomic Models
Two years later, Danish physicist, Niels Bohr,
proposed the Bohr Model of the atom
Atomic Models
Bohr’s Model
Electrons are
located certain
distances from
the nucleus
Each distance is
a certain quantity
of energy that the
electron can
have
Atomic Models
Electrons closest
to the nucleus
have the lowest
energy, while the
ones further
away are in
higher energy
levels
Atomic Models
The difference
between two
energy levels is
called a quantum
of energy.
Electrons can be
only in an energy
level, NOT
between levels.
Atomic Models
Electrons and Light (pg. 92)
Basic Information:
Light travels in waves
Each wave has a certain wavelength
(distance between two consecutive
peaks of a wave)
Atomic Models
Each wavelength has
a certain frequency (the # of waves that
pass through a specific point in one
second)
Atomic Model
Light and the Visible Electromagnetic Spectrum
 We can only see a small amount of the electromagnetic
spectrum
 Bohr found that if you pass a high voltage through a gas and
look at it through a prism, it will have a distinctive pattern of
colored lines called “line-emission spectrum”
Atomic Models
Bohr calculated that the line-emission
spectrum corresponded to energy levels in
the atom (which give off a different
wavelength of visible light)
The colors are caused when an electron falls
from its excited state to its ground state and
gives off energy in the form of light! (pg. 94)
Atomic Models
Great video on line-emission spectrum
http://www.mhhe.com/physsci/chemistry/essen
tialchemistry/flash/linesp16.swf
Spectrum of a Fluorescent Light
Spectrum