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Dalton's Atomic Theory
It was in the early 1800s that John Dalton, an observer of weather and
discoverer of color blindness among other things, came up with his atomic
theory. Let's set the stage for Dalton's work. Less than twenty years earlier, in
the 1780's, Lavoisier ushered in a new chemical era by making careful
quantitative measurements which allowed the compositions of compounds to
be determined with accuracy. By 1799 enough data had been accumulated for
Proust to establish the Law of Constant Composition ( also called the Law of
Definite Proportions). In 1803 Dalton noted that oxygen and carbon combined
to make two compounds. Of course, each had its own particular weight ratio
of oxygen to carbon (1.33:1 and 2.66:1), but also, for the same amount of
carbon, one had exactly twice as much oxygen as the other. This led him to
propose the Law of Simple Multiple Proportions, which was later verified by
the Swedish chemist Berzelius. In an attempt to explain how and why
elements would combine with one another in fixed ratios and sometimes also
in multiples of those ratios, Dalton formulated his atomic theory.
The idea of atoms had been proposed much earlier. The ancient Greek
philosophers had talked about atoms, but Dalton's theory was different in that
it had the weight of careful chemical measurements behind it. It wasn't just a
philosophical statement that there are atoms because there must be atoms.
His atomic theory, stated that elements consisted of tiny particles called
atoms. He said that the reason an element is pure is because all atoms of an
element were identical and that in particular they had the same mass. He also
said that the reason elements differed from one another was that atoms of
each element were different from one another; in particular, they had
different masses. He also said that compounds consisted of atoms of different
elements combined together. Compounds are pure substances (remember
they cannot be separated into elements by phase changes) because the atoms
of different elements are bonded to one another somehow, perhaps by hooks,
and are not easily separated from one another. Compounds have constant
composition because they contain a fixed ratio of atoms and each atom has its
own characteristic weight, thus fixing the weight ratio of one element to the
other. In addition he said that chemical reactions involved the rearrangement
of combinations of those atoms. So that, briefly, is Dalton's theory. With
modifications, it has stood up pretty well to the criteria that we talked about
earlier. It did not convince everyone right away however. Although a number
of chemists were quickly convinced of the truth of the theory, it took about a
half century for the opposition to die down, or perhaps I should say die off.
http://dl.clackamas.edu/ch104-04/dalton's.htm
Plum pudding model
The plum pudding model of the atom by J. J. Thomson, who discovered the electron in
1897, was proposed in 1904 before the discovery of the atomic nucleus. In this model,
the atom is composed of electrons (which Thomson still called "corpuscles", though G.
J. Stoney had proposed that atoms of electricity be called electrons in 1894[1])
surrounded by a soup of positive charge to balance the electron's negative charge, like
negatively-charged "plums" surrounded by positively-charged "pudding". The electrons
(as we know them today) were thought to be positioned throughout the atom, but with
many structures possible for positioning multiple electrons, particularly rotating rings of
electrons (see below). Instead of a soup, the atom was also sometimes said to have had a
cloud of positive charge. The model was disproved by the 1909 gold foil experiment,
which was interpreted by Ernest Rutherford in 1911[2] to imply a very small nucleus of
the atom containing a very high positive charge (enough to balance about 100 electrons
in gold), thus leading to the Rutherford model of the atom. Finally, after Henry
Moseley's work showed in 1913 that the nuclear charge was very close to the atomic
number, Antonius Van den Broek suggested that atomic number is nuclear charge. This
work had culminated in the solar-system-like (but quantum-limited) Bohr model of the
atom in the same year, in which a nucleus containing an atomic number of positive
charge is surrounded by an equal number of electrons in orbital shells.
Thomson's model was compared (though not by Thomson) to a British treat called plum
pudding, hence the name. It has also been called the chocolate chip cookie model or
blueberry muffin model, but these mental pictures assume the particles as static, which
they were not for Thomson. In this model, the electrons were free to rotate within the
blob or cloud of positive substance. These orbits were stabilized in the model by the fact
that when an electron moved farther from the center of the positive cloud, it felt a larger
net positive inward force, because there was more material of opposite charge, inside its
orbit (see Gauss's law). In Thomson's model, electrons were free to rotate in rings which
were further stabilized by interactions between the electrons, and spectra were to be
accounted for by energy differences of different ring orbits. Thomson attempted to
make his model account for some of the major spectral lines known for some elements,
but was not notably successful at this. Still, Thomson's model (along with a similar
Saturnian ring model for atomic electrons, put forward also in 1904 by Nagaoka after
James C. Maxwell's model of Saturn's rings), were earlier harbingers of the later and
more successful solar-system-like Bohr model of the atom.
http://en.wikipedia.org/wiki/Plum_pudding_model
Rutherford-Bohr Model
In atomic physics, the Bohr model, devised by Niels Bohr, depicts the atom as a small,
positively charged nucleus surrounded by electrons that travel in circular orbits around
the nucleus—similar in structure to the solar system, but with electrostatic forces
providing attraction, rather than gravity. This was an improvement on the earlier cubic
model (1902), the plum-pudding model (1904), the Saturnian model (1904), and the
Rutherford model (1911). Since the Bohr model is a quantum physics-based
modification of the Rutherford model, many sources combine the two, referring to the
Rutherford–Bohr model. Introduced by Niels Bohr in 1913, the model's key success
lay in explaining the Rydberg formula for the spectral emission lines of atomic
hydrogen. While the Rydberg formula had been known experimentally, it did not gain a
theoretical underpinning until the Bohr model was introduced. Not only did the Bohr
model explain the reason for the structure of the Rydberg formula, it also provided a
justification for its empirical results in terms of fundamental physical constants.
Bohr used to introduce his attempts to explain clearly the principles of the quantum
theory of the atom with an historical sketch, beginning invariably with the nuclear
model proposed by Rutherford. That was sound pedagogy but bad history. The
Rutherford-Bohr atom stands in the middle of a line of work initiated by J.J. Thomson
and concluded by the invention of quantum mechanics. Thompson's program derived its
inspiration from the peculiar emphasis on models characteristic of British physics of the
19th century. Rutherford's atom was a late product of the goals and conceptions of
Victorian science. Bohr's modifications, although ultimately fatal to Thomson's
program, initially gave further impetus to it. In the early 1920s the most promising
approach to an adequate theory of the atom appeared to be the literal and detailed
elaboration of the classical mechanics of multiply periodic orbits. The approach
succeeded, demonstrating in an unexpected way the force of an argument often
advanced by Thomson: because a mechanical model is richer in implications than the
considerations for which it was advanced, it can suggest new directions of research that
may lead to important discoveries