Download Key Historical Developments and Experiments that reveal Atoms

Key Historical Developments and Experiments that reveal Atoms and
Atomic Structure
Antiquity ~ 400 B.C.
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Greek philosophers Leucippus and Democritus asserted that all material things are
composed of extremely tiny, indivisible particles called atoms. “Nothing exists except
atoms and empty space. Everything else is opinion”. The atomic theory was
rejected by Aristotle, and thus by almost everyone else for the next two millennia.
Late 1700's through early 1800's: The idea of the atom is (re)developed.
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Antoine Lavoisier (in 1789) formulates and states the Law of Conservation of
Mass. He demonstrated this via some of the first quantitative experiments of the
time. By carefully weighing the reactants and products in chemical reactions, he
showed that the total mass of matter remains constant during the reaction (mass
before reaction equals mass after reaction). Lavoisier also performs a quantitative
study of combustion and elucidates the role of oxygen in this process.
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Joseph Louis Proust (in 1799) studies the composition of compounds and develops
the Law of Definite Proportions. This law states that if a compound is broken
down into its constituent elements, the masses of the elements will always occur in
fixed proportions, regardless of the quantity or source of the original substance.
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John Dalton (in 1805) proposes his Atomic Theory to explain the results of Lavoisier
and Proust, as well as his own experiment work and that of many other scientists.
Dalton’s Atomic Theory
a. Elements consist of tiny, indivisible particles called atoms.
b. All the atoms of a given element are identical, however, the atoms of different
elements differ in some fundamental way.
c. Compounds form when atoms of different elements combine in simple whole
number ratios. Thus, a given compound always contains the same relative
number and types of atoms.
d. During a chemical reaction, atoms are neither created nor destroyed. Instead,
reactions involve the reorganization of the atoms – a change in the way they are
grouped together. The atoms themselves are unaltered.
This was the first truly scientific theory of the atom, since Dalton reached his
conclusions by experimentation and examination of the results in an empirical
manner.
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Dalton proposes the Law of Multiple Proportions to help support his theory: When
several compounds can be formed from the same elements, for a fixed mass of one
element, the ratio of the masses of the other element will occur as small whole
number values. Experiments demonstrate that this law holds true.
Late 1800's through early 1900's: The internal structure of atoms is discovered.
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J.J. Thomson (in 1896) performs experiments with Cathode Ray Tubes (CRT) and
discovers the electron. The figure below is a simple schematic diagram of a CRT.
The original device used by Thomson was about one meter in length, and was made
entirely by hand.
Anodes / Collimators
Cathode
Electrical Plates
The entire glass tube was evacuated down to as low a vacuum as could be produced
(at that time), then sealed. Two electrical plates were placed about midway in the
CRT (connected to a powerful electric battery), through which the cathode rays were
passed. Thomson also used magnets, which were placed on either side of the
straight portion of the tube just to the right of the electrical plates. This allowed him
to use either electrical or magnetic fields (or a combination of both) to cause the
cathode rays to bend. Thomson discovered:
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The rays originate from the cathode (negative electrode) and travel to the anode
(positive electrode).
The rays consist of a stream of particles which have mass.
The particles in the rays have negative charge.
The same particles are generated no matter what material is used for the
cathode.
The charge to mass ratio of the particles, e/m = -1.76 x 108 C/g.
Thomson called these particles “corpuscles”. We now call these particles electrons.
They are present in all atoms.
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Thomson (in 1904) then proposes the Plum Pudding Model of the atom: An atom
consists of diffuse positive charge with negatively charged electrons imbedded in it,
like raisins in a pudding.
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Robert Millikan (in 1909) measures the charge on one electron. His experiment is
called the Millikan Oil Drop Experiment. A figure of the apparatus he used is
shown below.
Charged oil droplets were allowed to fall between two electrical plates that supply a
uniform electric field. The droplets were generated by an atomizer and a negative
charge was applied by irradiating them with X-rays. The potential of the electric field
was adjusted so that instead of falling, a drop would be suspended between the
plates in mid-air. The gravitational and electric forces on the charged droplets are
thus balanced.
When a drop is suspended, its weight m•g is exactly equal to the electric force
applied q•E. The values of E, the applied electric field, m the mass of a drop, and g,
the acceleration due to gravity, are all known values. So Millikan was able to solve
for q, the charge on the drop:
Repeating the experiment for many droplets, he discovered that the values
measured were always multiples of the same number. This number was determined
to be the charge on a single electron: 1.602 × 10−19 Coulombs.
Finally, using this value and the e/m ratio determined by Thomson, Millikan was able
to calculate the mass of an electron (= 9.11 x 10-31 kg).
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Radioactivity is discovered. Three types of radiation are identified:
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α rays (alpha rays) = tiny, heavy particles with a positive charge and high velocity
β rays (beta rays) = very high speed electrons
γ rays (gamma rays) = no mass, no charge, similar to x-rays but shorter
wavelength
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Ernest Rutherford (in 1909) discovers the nucleus of the atom and demonstrates that
the Plum Pudding Model is incorrect. His experiment is called the Gold Foil
Experiment (or Geiger-Marsden experiment).
A schematic diagram of this
experiment is shown below.
Rutherford and his team (Geiger and Marsden) bombarded a number of different thin
metal foils with α rays – tiny, heavy particles with a positive charge. Particles
passing through the foil would strike a screen coated with zinc sulfide, and appear as
tiny flashes of light.
Given the very high mass and momentum of the α particles, the expectation was that
the particles would pass through the foil and be scattered by tiny angles at most.
However, to their amazement, a few (~1 in 8000 particles) were deflected by large
angles (greater than 90°), an observation completely at odds with the predictions of
the Plum Pudding Model.
Based on his analysis of these scattering results, Rutherford proposes a revised
model of the atom called the Nuclear Model of the atom. In this model, all the
positive charge and most of the atomic mass is concentrated at the center of the
atom – in the nucleus – while electrons “orbit the nucleus, like planets orbit the sun.
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Rutherford (in 1918) was able to experimentally identify particles with positive charge
in the nucleus, which he called protons. But although he could explain the charge of
atomic nuclei with the right number of protons, the mass of an atom’s nucleus was
always larger than the sum of its protons. Therefore he postulated the existence of a
neutral particle with a mass nearly the same as the proton which, when added to the
protons in the nucleus, would give the right mass. Rutherford called this hypothetical
particle the neutron. Much later, James Chadwick (in 1930) was able to detect the
neutron experimentally.