Download Goal 4.01

Goal 5.01
Atomic Theory
Basic Laws of Chemistry
Law of conservation of mass: mass is neither
created or destroyed during chemical reactions
or physical changes
Law of definite proportions: compounds contain
the same proportions of elements by mass
regardless of the size of the sample
Law of multiple proportions: different compounds
made from the same elements result from the
combination of different whole numbers of atoms
Dalton
Dalton turned Democritus' idea of a
indivisible particle of matter into a theory
that could be tested.
He called it the atomic theory. It explained
the laws of chemistry that all atoms
followed.
Not all aspect’s of Dalton’s theory have
been proven correct.
Atomic Theory
All elements are composed of atoms that cannot
be divided.
Atoms of the same element have the same
mass and other properties.
Atoms cannot be created or destroyed.
Compounds contain the atoms of more than one
element.
In chemical reactions, atoms of the same
element always combine, separate, or rearrange
in the same way.
Where He Went Wrong
Parts of Dalton’s theory have been shown
to be untrue…
Atoms can be divided in to smaller parts.
Atoms of the same element can have
different masses.
Thomson’s
Experimental Setup
Thomson preformed an experiment using a
cathode ray tube.
It is a sealed tube that contains a gas at a low
pressure.
Current is passed through the gas from the
cathode (+) to anode (-) sides of the tube.
When this is done, a glowing beam of what looks
like light is created.
This glowing beam is called the cathode ray.
Thomson’s
Experimental Results
When a paddle wheel was placed in the
cathode ray, the wheel turned and moved.
When a negative magnetic field was
placed near the cathode ray, the ray bent
away from the charge.
When a positive magnetic field was placed
near the cathode ray, the ray bent toward
the charge.
Thomson’s
Conclusion
Because the paddle wheel turned, the beam
must actually be particles.
The mass of these particles was smaller than an
atom’s mass and therefore the particle must be
subatomic.
Because of how it was deflected by charges, the
particles must be negative.
Thompson named these negative, subatomic
particles electrons.
Inferences
This led to other inferences about atomic
structure.
Because atoms are neutral, there must be
a positive charge to balance the negatives.
Because electrons weigh so little their
must be something else to account for the
weight.
Rutherford’s
Experimental Setup
A thin piece of gold foil, full of tiny gold atoms,
was bombarded with fast moving alpha particles.
Alpha particles are a positive type of radiation
with four times the mass of a hydrogen.
It was thought that all alpha particles would pass
through the gold foil because nothing would be
big or strong enough to deflect them.
Rutherford’s
Experimental Results
BUT… 1 in every 8,000 alpha particles
were deflected back.
This deflection was caused by the
powerful repulsive force of a positively
charged space within the atom and the
density that the space had.
Rutherford’s
Conclusion
Rutherford concluded that the space was small
due to how few alpha particles were deflected.
Rutherford named the small, positively charged
space a nucleus and the space surrounding it an
electron cloud.
He named the particles inside the nucleus that
gave it it’s positive charge protons.
This positively charged protons were what was
canceling out the negative charge of the
electrons.
Bohr’s
Experimental Setup
Bohr made electrons absorb energy and
jump up in orbit. When this occurs, the
atom is said to be excited.
As the electron falls from the excited state
to the ground state, energy is given back off
in the form of a photon of radiation.
The photon of energy was then sent
through a prism where it formed a
spectrum.
Bohr’s
Experimental Results
The spectrum that was produced was
a line emission spectrum.
This indicated that only fixed amounts
of energy were being
released/absorbed as electrons moved
between orbits.
Bohr’s
Conclusion
Bohr discovered that electrons don’t just move
randomly around a nucleus because if they did
the spectrum would have been continuous.
Instead the electrons follow distinct paths called
orbits and these orbits are at set distances from
one another.
Bohr used the wavelengths of the lines in the
spectrum to calculate the exact energy that the
electrons have at each level.
Electron Probability
It is now known that electrons don’t have
to follow these paths.
Electrons are just likely to follow paths
called orbits.
They have a probability to.