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					Development of Atomic Theory
"Atomic model" redirects here. For the unrelated term in mathematical logic, see Atomic model
(mathematical logic).
      This article focuses on the historical models of the atom. For a history of the study of how
      atoms combine to form molecules, see History of the molecule.

In chemistry and physics, atomic theory is a theory of the nature of matter, which states that
matter is composed of discrete units called atoms, as opposed to the obsolete notion that matter
could be divided into any arbitrarily small quantity. It began as a philosophical concept in
ancient Greece and India and entered the scientific mainstream in the early 19th century when
discoveries in the field of chemistry showed that matter did indeed behave as if it were made up
of particles.

The word "atom" (from the Greek atomos, "indivisible") was applied to the basic particle that
constituted a chemical element, because the chemists of the era believed that these were the
fundamental particles of matter. However, around the turn of the 20th century, through various
experiments with electromagnetism and radioactivity, physicists discovered that the so-called
"indivisible atom" was actually a conglomerate of various subatomic particles (chiefly, electrons,
protons and neutrons) which can exist separately from each other. In fact, in certain extreme
environments such as neutron stars, extreme temperature and pressure prevents atoms from
existing at all. Since atoms were found to be actually divisible, physicists later invented the term
"elementary particles" to describe indivisible particles. The field of science which studies
subatomic particles is particle physics, and it is in this field that physicists hope to discover the
true fundamental nature of matter.

      What experimental measurements were available to Dalton?

 John Dalton was able to supply experimental results to forcefully revive the idea
                                   of the atom.

  He was influenced by the experiments of two Frenchmen, Antoine Lavoisier and
                              Joseph Louis Proust.



Antoine Lavoisier (1743-1794) - Formulated the Law of Conservation of Matter: "Matter is
neither gained nor lost during a chemical reaction." He did this by weighing materials before
and after reactions. For example, the weights of the mercury and oxygen formed by
decomposition of mercuric oxide were compared with the initial weight of the mercuric oxide.

Joseph Louis Proust (1754-1826) - Formulated the Law of Constant Porportions: "In a
compound, the contsitutne elements are always present in a definite proportion by weight." Like
Lavoisier, Proust also conducted quantitative experiments. He showed that regardless of how
copper carbonate was prepared in the laboratory, or how it was isolated from nature, it always
contained the same proportions of copper, oxygen and carbon - 5:4:1 parts by weight.

Read about Proust's research on copper in his own words.

Not all his contemporaries agreed with Proust's conclusions. Berthollet was able to combine
different quantities of copper and tin to produce what seemed to him to be compounds of varying
composition. What is the difference between the combination of carbon and oxygen in carbon
dioxide, and the combination of copper and tin when they are heated together?

John Dalton (1766-1844)- Formulated the Law of Multiple Proportions : "In the formation of
two or more compounds from the same elements, the weights of one element that combine with a
fixed weight of a second element are in a ratio of small whole numbers (integers) such as 2 to 1,
3 to 1, 3 to 2, or 4 to 3." He had made a quantitative study of different compounds made from the
same elements, such as carbon monoxide and carbon dioxide. He found that the weight ratio of
carbon to oxygen in carbon monoxide was 3:4, and the weight ratio of carbon to oxygen in
carbon dioxide was 3:8.

Read a short article about the 'Chemical atom in early 19th century chemistry' which
describes the period including the work of the three scientists mentioned above.

Read John Dalton's own words as he discusses the opinions of some of his contemporaries and
gives his own ideas about how elements combine to form compounds



N       ear the end of the 18th century, two laws about chemical reactions emerged without
        referring to the notion of an atomic theory. The first was the law of conservation of
        mass, formulated by Antoine Lavoisier in 1789, which states that the total mass in a
chemical reaction remains constant (that is, the reactants have the same mass as the products).
The second was the law of definite proportions. First proven by the French chemist Joseph Louis
Proust in 1799, this law states that if a compound is broken down into its constituent elements,
then the masses of the constituents will always have the same proportions, regardless of the
quantity or source of the original substance. Proust had synthesized copper carbonate through
numerous methods and found that in each case the ingredients combined in the same proportions
as they were produced when he broke down natural copper carbonate.
Various atoms and molecules as depicted in John Dalton's A New System of Chemical
Philosophy (1808).

In the early years of the 19th century, John Dalton developed an atomic theory in which he
proposed that each chemical element is composed of atoms of a single, unique type, and that
though they are both immutable and indestructible, they can combine to form more complex
structures (chemical compounds). The conservation of mass suggested to Dalton that the atoms
of matter are indestructible. His theory allowed him to explain various new discoveries in
chemistry that he and his contemporaries made. This marked the first truly scientific theory of
the atom, since Dalton reached his conclusions by experimentation and examination of the
results in an empirical fashion. It is unclear to what extent his atomic theory might have been
inspired by earlier such theories.

Dalton studied and expanded upon Proust's work to develop the law of multiple proportions: if
two elements form more than one compound between them, then the ratios of the masses of the
second element which combine with a fixed mass of the first element will be ratios of small
integers. One pair of reactions Dalton studied involved the combinations of "nitrous air", or what
we now call nitric oxide (NO), and oxygen (O2). Under certain conditions, these gases formed an
unknown product at a certain combining ratio (now known to be nitrogen dioxide (NO2)), but
when he repeated the reaction under other conditions, exactly twice the amount of nitric oxide (a
ratio of 1:2—small integers) reacted completely with oxygen to form a different product—now
known as dinitrogen trioxide (N2O3).

2NO + O2 → 2NO2

4NO + O2 → 2N2O3

Dalton also believed atomic theory could explain why water absorbed different gases in different
proportions: for example, he found that water absorbed carbon dioxide far better than it absorbed
nitrogen. Dalton hypothesized this was due to the differences in mass and complexity of the
gases' respective particles. Indeed, carbon dioxide molecules (CO2) are heavier and larger than
nitrogen molecules (N2).
In 1803 Dalton orally presented his first list of relative atomic weights for a number of
substances. This paper was published in 1805, but he did not discuss there exactly how he
obtained these figures. The method was first revealed in 1807 by his acquaintance Thomas
Thomson, in the third edition of Thomson's textbook, A System of Chemistry. Finally, Dalton
published a full account in his own textbook, A New System of Chemical Philosophy, 1808 and
1810.

Dalton estimated the atomic weights according to the mass ratios in which they combined, with
hydrogen being the basic unit. However, Dalton did not conceive that with some elements atoms
exist in molecules – e.g. pure oxygen exists as O2. He also mistakenly believed that the simplest
compound between any two elements is always one atom of each (so he thought water was HO,
not H2O). This, in addition to the crudity of his equipment, resulted in his table being highly
flawed. For instance, he believed oxygen atoms were 5.5 times heavier than hydrogen atoms,
because in water he measured 5.5 grams of oxygen for every 1 gram of hydrogen and believed
the formula for water was HO (an oxygen atom is actually 16 times heavier than a hydrogen
atom).

The flaw in Dalton's theory was corrected in 1811 by Amedeo Avogadro. Avogadro had
proposed that equal volumes of any two gases, at equal temperature and pressure, contain equal
numbers of molecules (in other words, the mass of a gas's particles does not affect its volume).
Avogadro's law allowed him to deduce the diatomic nature of numerous gases by studying the
volumes at which they reacted. For instance: since two liters of hydrogen will react with just one
liter of oxygen to produce two liters of water vapor (at constant pressure and temperature), it
meant a single oxygen molecule splits in two in order to form two particles of water. Thus,
Avogadro was able to offer more accurate estimates of the atomic mass of oxygen and various
other elements, and firmly established the distinction between molecules and atoms.

In 1815 the English chemist William Prout observed that the atomic weights that had been
measured for the elements known at that time appeared to be whole multiples of the atomic
weight of hydrogen. Prout hypothesized that the hydrogen atom was the only truly fundamental
object, and that the atoms of other elements were actually groupings of various numbers of
hydrogen atoms. Prout's hypothesis was confirmed in essence by Ernest Rutherford a century
later.

In 1827, the British botanist Robert Brown observed that pollen particles floating in water
constantly jiggled about for no apparent reason. In 1905, Albert Einstein theorized that this
Brownian motion was caused by the water molecules continuously knocking the grains about,
and developed a hypothetical mathematical model to describe it. This model was validated
experimentally in 1908 by French physicist Jean Perrin, thus providing additional validation for
particle theory (and by extension atomic theory).

Discovery of subatomic particles
Thomson's Crookes tube in which he observed the deflection of cathode rays by an electric field.
The purple line represents the deflected electron stream.

Atoms were thought to be the smallest possible division of matter until 1897 when J.J. Thomson
discovered the electron through his work on cathode rays. A Crookes tube is a sealed glass
container in which two electrodes are separated by a vacuum. When a voltage is applied across
the electrodes, cathode rays are generated, creating a glowing patch where they strike the glass at
the opposite end of the tube. Through experimentation, Thomson discovered that the rays could
be deflected by an electric field (in addition to magnetic fields, which was already known). He
concluded that these rays, rather than being waves, were composed of negatively charged
particles he called "corpuscles" (they would later be renamed electrons by other scientists).

Thomson believed that the corpuscles emerged from the very atoms of the electrode. He thus
concluded that atoms were divisible, and that the corpuscles were their building blocks. To
explain the overall neutral charge of the atom, he proposed that the corpuscles were distributed in
a uniform sea or cloud of positive charge; this was the plum pudding model as the electrons were
embedded in the positive charge like plums in a plum pudding.




Discovery of the nucleus




The gold foil experiment
Top: Expected results: alpha particles passing through the plum pudding model of the atom with
negligible deflection.
Bottom: Observed results: a small portion of the particles were deflected, indicating a small,
concentrated positive charge.
Thomson's plum pudding model was disproved in 1909 by one of his former students, Ernest
Rutherford, who discovered that most of the mass and positive charge of an atom is concentrated
in a very small fraction of its volume, which he assumed to be at the very center.

In the gold foil experiment, Hans Geiger and Ernest Marsden (colleagues of Rutherford working
at his behest) shot alpha particles through a thin sheet of gold, striking a fluorescent screen that
surrounded the sheet. Given the very small mass of the electrons, the high momentum of the
alpha particles and the unconcentrated distribution of positive charge of the plum pudding model,
the experimenters expected all the alpha particles to either pass through without significant
deflection or be absorbed. To their astonishment, a small fraction of the alpha particles
experienced heavy deflection.

This led Rutherford to propose a model of the atom (the planetary model or Rutherford model) to
explain the experimental results. In this model, the atom was made up of a nucleus of
approximately 10-15 m in diameter, surrounded by an electron cloud of approximately 10-10 m in
diameter. The pointlike electrons orbited in the space around the massive, compact nucleus like
planets orbiting the Sun.

Following this discovery, the study of the atom split into two distinct fields, nuclear physics,
which studies the properties and structure of the nucleus of atoms, and atomic physics, which
examines the properties of the electrons surrounding the nucleus.



First steps towards a quantum physical model of the atom

The planetary model of the atom had two significant shortcomings. The first is that, unlike the
planets orbiting the sun, electrons are charged particles. An accelerating electric charge is known
to emit electromagnetic waves according to the Larmor formula in classical electromagnetism;
an orbiting charge would steadily lose energy and spiral towards the nucleus, colliding with it in
a small fraction of a second. The second problem was that the planetary model could not explain
the highly peaked emission and absorption spectra of atoms that were observed.




The Bohr model of the atom

Quantum theory revolutionized physics at the beginning of the 20th century, when Max Planck
and Albert Einstein postulated that light energy is emitted or absorbed in discrete amounts
known as quanta (singular, quantum). In 1913, Niels Bohr incorporated this idea into his Bohr
model of the atom, in which the electrons could only orbit the nucleus in particular circular orbits
with fixed angular momentum and energy, their distances from the nucleus (i.e., their radii)
being proportional to their respective energies. Under this model electrons could not spiral into
the nucleus because they could not lose energy in a continuous manner; instead, they could only
make instantaneous "quantum leaps" between the fixed energy levels. When this occurred, light
was emitted or absorbed at a frequency proportional to the change in energy (hence the
absorption and emission of light in discrete spectra).

Bohr's model was not perfect. It could only predict the spectral lines of hydrogen; it couldn't
predict those of multielectron atoms. Worse still, as spectrographic technology improved,
additional spectral lines in hydrogen were observed which Bohr's model couldn't explain. In
1916, Arnold Sommerfeld added elliptical orbits to the Bohr model to explain the extra emission
lines, but this made the model very difficult to use, and it still couldn't explain more complex
atoms.




Discovery of isotopes
While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick
Soddy discovered that there appeared to be more than one element at each position on the
periodic table. The term isotope was coined by Margaret Todd as a suitable name for these
elements.

That same year, J.J. Thomson conducted an experiment in which he channeled a stream of neon
ions through magnetic and electric fields, striking a photographic plate at the other end. He
observed two glowing patches on the plate, which suggested two different deflection trajectories.
Thomson concluded this was because some of the neon ions had a different mass. The nature of
this differing mass would later be explained by the discovery of neutrons in 1932.

Discovery of nuclear particles

In 1918, Rutherford bombarded nitrogen gas with alpha particles and observed hydrogen nuclei
being emitted from the gas. Rutherford concluded that the hydrogen nuclei emerged from the
nuclei of the nitrogen atoms themselves (in effect, he split the atom). He later found that the
positive charge of any atom could always be equated to that of an integer number of hydrogen
nuclei. This, coupled with the facts that hydrogen was the lightest element known and that the
atomic mass of every other element was roughly equivalent to an integer number of hydrogen
atoms, led him to conclude hydrogen nuclei were singular particles and a basic constituent of all
atomic nuclei: the proton. Further experimentation by Rutherford found that the nuclear mass of
most atoms exceeded that of the protons it possessed; he speculated that this surplus mass was
composed of hitherto unknown neutrally charged particles, which were tentatively dubbed
"neutrons".

In 1928, Walter Bothe observed that beryllium emitted a highly penetrating, electrically neutral
radiation when bombarded with alpha particles. It was later discovered that this radiation could
knock hydrogen atoms out of paraffin wax. Initially it was thought to be high-energy gamma
radiation, since gamma radiation had a similar effect on electrons in metals, but James Chadwick
found that the ionisation effect was too strong for it to be due to electromagnetic radiation. In
1932, he exposed various elements, such as hydrogen and nitrogen, to the mysterious "beryllium
radiation", and by measuring the energies of the recoiling charged particles, he deduced that the
radiation was actually composed of electrically neutral particles with a mass similar to that of a
proton. For his discovery of the neutron, Chadwick received the Nobel Prize in 1935.

Quantum physical models of the atom
In 1924, Louis de Broglie proposed that all moving particles–particularly subatomic particles
such as electrons–exhibit a degree of wave-like behavior. Erwin Schrödinger, fascinated by this
idea, explored whether or not the movement of an electron in an atom could be better explained
as a wave rather than as a particle. Schrödinger's equation, published in 1926, describes an
electron as a wavefunction instead of as a point particle. This approach elegantly predicted many
of the spectral phenomena that Bohr's model failed to explain. Although this concept was
mathematically convenient, it was difficult to visualize, and faced opposition. One of its critics,
Max Born, proposed instead that Schrödinger's wavefunction described not the electron but
rather all its possible states, and thus could be used to calculate the probability of finding an
electron at any given location around the nucleus.




The five filled atomic orbitals of a neon atom, separated and arranged in order of increasing
energy from left to right, with the last three orbitals being equal in energy. Each orbital holds up
to two electrons, which exist for most of the time in the zones represented by the colored
bubbles. Each electron is equally in both orbital zones, shown here by color only to highlight the
different wave phase.

A consequence of describing electrons as waveforms is that it is mathematically impossible to
simultaneously derive the position and momentum of an electron; this became known as the
Heisenberg uncertainty principle. This invalidated Bohr's model, with its neat, clearly defined
circular orbits. The modern model of the atom describes the positions of electrons in an atom in
terms of probabilities. An electron can potentially be found at any distance from the nucleus,
but—depending on its energy level—tends to exist more frequently in certain regions around the
nucleus than others; this pattern is referred to as its atomic orbital.
1803 - John Dalton - Atomic Theory




    1.   Matter is made up of indivisible atoms.
    2.   All atoms of an element are identical.
    3.   Atoms are neither created nor destroyed.
    4.   Atoms of different elements have different weights and chemical properties.
    5.   Atoms of different elements combine in simple whole numbers to form compounds.

                                 Dalton's atomic theory
 Here is a summary of Dalton's theory.

 1. Elements are composed of tiny, separate, indivisible and indestructible particles. These
 particles, called atoms, maintain their identity when the element undergoes physical or
 chemical change.

 2. All atoms of the same element are identical and different from the atoms of every other
 element.

 3. Atoms combine in simple whole number ratios to form compounds.

 4. Atoms of the same elements can combine in different ratios to form more than one
 compound.

 A simple discussion of this theory and its background can be found here.

 Berzelius, contributed significantly to the development of atomic theory. About 1807 he
 performed a great number of analyses of chemical compounds, and showed so many examples
 of the law of definite proportions that it could no longer be doubted. He also set about
 determining atomic weights and his first table, published in 1828, compared favourably with
 today's accepted values.

 Whereas Dalton represented atoms of elements by circles containing a letter or symbol,
 Berzelius chose to omit the circle and just use an initial letter of the Latin name of the element
 (or two letters if more than one element began with the same letter). This led to the system we
 now use for writing formulae of elements and compounds and writing chemical equations.



 Now, 200 years later, how would you modify Dalton's theory?
Think of sub-atomic particles - protons, neutrons, electrons...

spontaneous fission of radioactive atoms, nuclear fission and fusion,

production of radioactive isotopes in an atomic pile... So the first part of the theory is no
longer accepted.

Atoms of isotopes of an element are not identical. So the second part of the theory is no longer
accepted.

The Modern Atomic Theory

   John Dalton, an English chemist, might be called "the father of the modern atomic theory."
During the early 1800's, Dalton's interests in Meteorology and gases lead him to read the
works of Antoine Lavoisier and Joseph Proust. Lavoisier had stated the law of conservation of
mass, that the mass of materials before a chemical reaction takes place is exactly equal to the
mass of the materials after the reaction is completed. Proust had observed the law of definite
proportions, stating that the proportion by mass of the elements in a given compound is always
the same. Dalton felt that the findings of these men gave strong support to the idea of atoms.
He formulated an atomic theory that would include the observations found by Lavoisier and
Proust.

                                   Dalton's Atomic Theory
          1) All elements are composed of atoms, which are indivisible and
          indestructible particles.
          2) All atoms of the same element are exactly alike; in particular, they all
          have the same mass.
          3) All atoms of different elements are different; in particular, they have
          different masses.
          4) Compounds are formed by the joining of atoms of two or more
          elements. In any compound , the atoms of the different elements in the
          compound are joined in a definite whole-number ratio, such as 1 to 1, 2 to
          1, 3 to 2, etc.

   Much has happened since the time of Dalton, which has made it necessary to update his
atomic theory. We currently believe that all elements are composed of atoms, but we know
that those atoms are not indestructible. Atoms are split in nuclear reactions, and they are made
up of even smaller particles. We also know that atoms of the same element can have different
masses, when they represent different isotopes. Despite these differences, much of Dalton's
atomic theory remains useful to this day.
Antoine Lavoisier (1743-1794) - Formulated the Law of Conservation of Matter: "Matter is
neither gained nor lost during a chemical reaction." He did this by weighing materials before
and after reactions. For example, the weights of the mercury and oxygen formed by
decomposition of mercuric oxide were compared with the initial weight of the mercuric oxide.

Joseph Louis Proust (1754-1826) - Formulated the Law of Constant Porportions: "In a
compound, the contsitutne elements are always present in a definite proportion by weight."
Like Lavoisier, Proust also conducted quantitative experiments. He showed that regardless of
how copper carbonate was prepared in the laboratory, or how it was isolated from nature, it
always contained the same proportions of copper, oxygen and carbon - 5:4:1 parts by weight.

Read about Proust's research on copper in his own words.

Not all his contemporaries agreed with Proust's conclusions. Berthollet was able to combine
different quantities of copper and tin to produce what seemed to him to be compounds of
varying composition. What is the difference between the combination of carbon and oxygen in
carbon dioxide, and the combination of copper and tin when they are heated together?

John Dalton (1766-1844)- Formulated the Law of Multiple Proportions : "In the formation of
two or more compounds from the same elements, the weights of one element that combine with
a fixed weight of a second element are in a ratio of small whole numbers (integers) such as 2
to 1, 3 to 1, 3 to 2, or 4 to 3." He had made a quantitative study of different compounds made
from the same elements, such as carbon monoxide and carbon dioxide. He found that the
weight ratio of carbon to oxygen in carbon monoxide was 3:4, and the weight ratio of carbon
to oxygen in carbon dioxide was 3:8.

Read a short article about the 'Chemical atom in early 19th century chemistry' which
describes the period including the work of the three scientists mentioned above.

Read John Dalton's own words as he discusses the opinions of some of his contemporaries
and gives his own ideas about how elements combine to form compounds.

From Dalton to the Periodic Table

Modern atomic theory begins with the work of John Dalton, published in 1808. He held that
all the atoms of an element are of exactly the same size and weight (see atomic weight) and
are in these two respects unlike the atoms of any other element. He stated that atoms of the
elements unite chemically in simple numerical ratios to form compounds. The best evidence
for his theory was the experimentally verified law of simple multiple proportions, which gives
a relation between the weights of two elements that combine to form different compounds.

Evidence for Dalton's theory also came from Michael Faraday's law of electrolysis. A major
development was the periodic table, devised simultaneously by Dmitri Mendeleev and J. L.
Meyer, which arranged atoms of different elements in order of increasing atomic weight so
that elements with similar chemical properties fell into groups. By the end of the 19th cent. it
was generally accepted that matter is composed of atoms that combine to form molecules.

Discovery of the Atom's Structure
In 1911, Ernest Rutherford developed the first coherent explanation of the structure of an
atom. Using alpha particles emitted by radioactive atoms, he showed that the atom consists of
a central, positively charged core, the nucleus, and negatively charged particles called
electrons that orbit the nucleus. There was one serious obstacle to acceptance of the nuclear
atom, however. According to classical theory, as the electrons orbit about the nucleus, they are
continuously being accelerated (see acceleration), and all accelerated charges radiate
electromagnetic energy. Thus, they should lose their energy and spiral into the nucleus.

This difficulty was solved by Niels Bohr (1913), who applied the quantum theory developed
by Max Planck and Albert Einstein to the problem of atomic structure. Bohr proposed that
electrons could circle a nucleus without radiating energy only in orbits for which their orbital
angular momentum was an integral multiple of Planck's constant h divided by 2π. The discrete
spectral lines (see spectrum) emitted by each element were produced by electrons dropping
from allowed orbits of higher energy to those of lower energy, the frequency of the photon of
light emitted being proportional to the energy difference between the orbits.

Around the same time, experiments on x-ray spectra (see X ray) by H. G. J. Moseley showed
that each nucleus was characterized by an atomic number, equal to the number of unit positive
charges associated with it. By rearranging the periodic table according to atomic number
rather than atomic weight, a more systematic arrangement was obtained. The development of
quantum mechanics during the 1920s resulted in a satisfactory explanation for all phenomena
related to the role of electrons in atoms and all aspects of their associated spectra. With the
discovery of the neutron in 1932 the modern picture of the atom was complete.

Contemporary Studies of the Atom

With many of the problems of individual atomic structure and behavior now solved, attention
has turned to both smaller and larger scales. On a smaller scale the atomic nucleus is being
studied in order to determine the details of its structure and to develop sources of energy from
nuclear fission and fusion (see nuclear energy), for the atom is not at all indivisible, as the
ancient philosophers thought, but can undergo a number of possible changes. On a larger scale
new discoveries about the behavior of large groups of atoms have been made (see solid-state
physics). The question of the basic nature of matter has been carried beyond the atom and now
centers on the nature of and relations between the hundreds of elementary particles that have
been discovered in addition to the proton, neutron, and electron. Some of these particles have
been used to make new types of exotic ―atoms‖ such as positronium (see antiparticle) and
muonium (see muon).




Atomic Theory:

The ancient philosopher, Heraclitus, maintained that everything is in a state of flux.
Nothing escapes change of some sort (it is impossible to step into the same river).
On the other hand, Parmenides argued that everything is what it is, so that it cannot
become what is not (change is impossible because a substance would have to
transition through nothing to become something else, which is a logical
contradiction). Thus, change is incompatible with being so that only the permanent
aspects of the Universe could be considered real.

An ingenious escape was proposed in the fifth century B.C. by Democritus. He
hypothesized that all matter is composed of tiny indestructible units, called atoms.
The atoms themselves remain unchanged, but move about in space to combine in
various ways to form all macroscopic objects. Early atomic theory stated that the
characteristics of an object are determined by the shape of its atoms. So, for
example, sweet things are made of smooth atoms, bitter things are made of sharp
atoms.

In this manner permanence and flux are reconciled and the field of atomic physics
was born. Although Democritus' ideas were to solve a philosophical dilemma, the
fact that there is some underlying, elemental substance to the Universe is a primary
driver in modern physics, the search for the ultimate subatomic particle.




It was John Dalton, in the early 1800's, who determined that each chemical element
is composed of a unique type of atom, and that the atoms differed by their masses.
He devised a system of chemical symbols and, having ascertained the relative
weights of atoms, arranged them into a table. In addition, he formulated the theory
that a chemical combination of different elements occurs in simple numerical ratios
by weight, which led to the development of the laws of definite and multiple
proportions.
He then determined that compounds are made of molecules, and that molecules are
composed of atoms in definite proportions. Thus, atoms determine the composition
of matter, and compounds can be broken down into their individual elements.

The first estimates for the sizes of atoms and the number of atoms per unit volume
where made by Joesph Loschmidt in 1865. Using the ideas of kinetic theory, the
idea that the properties of a gas are due to the motion of the atoms that compose it,
Loschmidt calculated the mean free path of an atom based on diffusion rates. His
result was that there are 6.022x1023 atoms per 12 grams of carbon. And that the
typical diameter of an atom is 10-8 centimeters.

By the 19th century, it was determined that atoms bind together to form substances
through electromagnetic forces. That atoms are very small and matter is mostly
empty space, but 'feels' solid because the atoms of your hand are repelled by the
electromagnetic forces between your atoms and an objects atoms (like a table).
1830 - Michael Faraday
     Set up a pair of metal plates sealed in a glass tube. The tube was filled with a gas, and
      the metal plates were connected to a series of batteries.

                      As the pressure of the gas decreased, the gas began to glow.
                      Julius Plucker (1858) noticed that only one end emitted light.
                           o He also changed the position of the patch of glass that glowed by
                              bringing
                              a magnet close to the tube.
                      Conclusion: The effect of the magnetic field as evidence that
                       whatever
                       produced this glow was electrically charged.
                      Cathode - metal plate connected to the negative end
                      Anode - metal plate connected to the positive end




1869 - Johannes Hittorf
     Found that when a solid object was placed between the cathode and anode, a shadow was
      cast on the end of the tube across from the cathode.
     Conclusion: Some beam or ray is given off by the cathode - subsequently called the
      tubes cathode-ray tubes.

1879 - William Crookes

                          Developed a better vacuum pump that allowed him to produce
                           cathode-ray
                           tubes with a smaller residual gas pressure.
                          Conclusion: Cathode r0ays are negatively charged by studying
                           deflection
                           of cathode rays by magnetic fields.



1897 - J.J. Thompson
                           Found that cathode rays could be deflected by an electric field
                           Showed that cathode "rays" were actually particles
                           Found the charge to mass ratio of the particles to be
                            approximately
                            108 Coulomb (C) per gram.
                           Same charge to mass ratio regardless of metal used for
                            cathode/anode
                            or gas used to fill the tube.
                           Conclusion: Particles were a universal component of matter.

                           Electron - (originally called corpuscles by Thompson) particles
                            given off by the cathode; fundamental unit of negative electricity
                           Raisin Pudding Model -
                                o Matter is electrically neutral and electrons are much lighter
                                   than atoms.
                                o Conclusion: There must be positively charged
 Raisin Pudding                    particles which also must carry the mass of the
     Model                         atom. Dalton's model is now incorrect because
                                   atoms are divisible.



1895 - William Conrad Roentgen
                     Discovered x-rays while using cathode-ray tubes. Found that x-rays
                      could pass
                      through solid objects.




1899 - Ernst Rutherford
             Studied absorption of radioactivity.

                     Alpha radiation - positive charge - absorbed by a few hundredths of a
                      cm or metal foil
                     Beta radiation - negative charge - could pass through 100x as much foil
                      before it was absorbed
                     Gamma rays - no charge - could penetrate several cm of lead
1907-1911 - Rutherford updated Thomson's Raisin Pudding Model of the
atom.

       Studied the deflection of alpha particles as they were
        targeted
        at thin gold foil sheets.
             o Most of the alpha particles penetrated straight
                through.
             o However few were deflected at slight angles.
             o Even fewer (only about 1 in 20,000) were
                deflected at
                angles over 90 .




       Conclusion: The positive charge and mass of an atom were concentrated in the
        center and only made up a small fraction of the total volume. He named this
        concentrated center the nucleus (Latin for little nut).
       Rutherford was also able to estimate the charge of an atom by studying the deflection of
        alpha particles. He found that the positive charge on the atom was approximately half of
        the atomic weight.

1908-1917 - Robert Millikan - Oil-drop experiment

                        J.J. Thomson had previously hypothesized that the mass of a single
                         electron
                         was at least 1000 times smaller than that of the smallest atom.
                        Millikan measured the charge on an electron with his oil-drop apparatus.
     An "atomizer" from a
      perfume bottle sprayed
      oil or water droplets into
      the sample chamber.
      Some of the droplets fell
      through the pinhole
      into an area between two
      plates (one positive
      and one negative). This
      middle chamber was
      ionized by x-rays.
       Particles that did not
      capture any electrons fell
      to the bottom plate
      due to gravity. Particles
      that did capture one
      or more electrons were
      attracted to the
      positive upper plate and
      either floated upward
      or fell more slowly.


     Conclusion: The charge on a drop was always a multiple of 1.59 x 10-19 Coulombs.
       He proved Thomson's hypothesis that the mass of an electron is at least 1000 times
      smaller than the smallest atom.




1913 - A. van den Broek
     Suggested that the positive charge on atoms should be compared to their atomic numbers,
      not their atomic weights.
          o At the time, atomic number (Z) only specified an element's location on the
              periodic table. Today, the atomic number is, by definition, the number of protons
              in an atom.

1914 - H. G. J. Moseley
                        Studied the frequencies of the x-rays given off by cathode-ray tubes
                         when electrons strike the anode.
                        Found that there was a relationship between the frequencies (v) of the
                         x-rays given off by the cathode-ray tube and the atomic number of the
                         metal used to form the anode:
       Conclusion: He argued that the frequencies of the x-rays should depend on the
        charge on the nucleus emitting these x-rays. Therefore, the atomic number was
        equal to the positive charge (charge on the nucleus) of an atom.



1920 - Rutherford proposed the name "proton" for the positively charged particles in the
nucleus of an atom. At the same time, he also proposed that the nucleus also contained
electrically neutral particles which accounted for the remaining mass of the atom. He called this
yet unknown particle the "neutron".

1932 - James Chadwick
                        Proved that neutrons, neutral particles in the nucleus that made up
                         approximately
                         half the mass of an atom, did exist.

                               Summary of Subatomic Particles
                                Particle Symbol Charge                  Mass
                                                    -
                                 Electron       e          -1      0.0005486 amu
                                                 +
                                 Proton         p          +1       1.007276 amu
                                 Neutron        n           0       1.008665 amu

                                      Atomic Rules
       The number of protons in the nucleus of an atom is equal to the atomic number (Z).
       In a neutral atom, the number of electrons is equal to the number of protons.
       The mass number (M) of an atom is equal to the sum of the number of
        protons and neutrons in the nucleus.
       The number of neutrons is equal to the difference between the mass number (M)
        and the atomic number (Z).




                Atomic number: protons (and electrons if neutral)

                Mass number: protons + neutrons (neutrons = mass number - atomic number)
    Ancient Greeks struggled to understand the nature of matter
Empedocles (around 490 to 444 BC) thought there were four original elements: Earth, Air, Fire,
Water. He thought everything else came about through their combination and/or separation by
the two opposite principles of Love and Strife.

Leucippus (around 460 to 420 BC) and Democritus (around 460 to 370 BC), supposedly a pupil
of Leucippus, are considered the founders of atomism. Leucippus regarded atoms as
imperceptible, individual particles that differ only in shape and position.

Plato (about 427 to 347 BC) in his work, the Timaeus, proposes a mathematical construction of
the elements - earth, air, fire, water. Each of these elements is said to consist of particles or
primary bodies. Each particle is a regular geometrical solid- the cube, tetrahedron, octahedron
and icosahedron. Each of these particles is composed of simple right triangles. The particles are
like the molecules of the theory; the triangles are its atoms.

Plato's beliefs as regards the universe were that the stars, planets, Sun and Moon move round the
Earth in crystalline spheres. The sphere of the Moon was closest to the Earth, then the sphere of
the Sun, then Mercury, Venus, Mars, Jupiter, Saturn and furthest away was the sphere of the
stars. He believed that the Moon shines by reflected sunlight.

Aristotle (384-322 BC) said that Earth was both the centre of the universe and one of the four
primordial elements. He saw the universe as a series of concentric spheres, with earth at the
centre, followed by water, air, fire. The harmonious relationships and interworkings of these
spheres could be heard as celestial music: the music of the spheres. Above fire was the Moon,
and this sphere delimited matter of a different kind. Beyond the Moon were spheres for the Sun,
the planets, and the stars, which were carried around the Earth in daily, complicated orbits. All
matter inside of the Moon’s orbit was different in kind from matter above the Moon.
Reminiscent of Plato’s ideas, Aristotle’s theory said that terrestrial matter decays and is
ephemeral, while celestial matter, the aether, is unchanging and eternal.




Rutherford Atomic Theory was first posited in 1911, many facets of it are still accepted by the
majority of the scientific community.

       History of
   1. In the early 1900s the predominant theory regarding how atoms were physically
      structured was called the Plum Pudding Model. The Plum Pudding Model had been
      theorized and established by J. J. Thomson. In 1911, physicist Ernest Rutherford used his
      experimental data from several different experiments to conclude that the basic atomic
      structure was very different than that proposed by the Plum Pudding Model. Rutherford's
      work eventually led to the coherent theory which we refer to today as Rutherford's
      Atomic Theory.
    Features

2. The basic difference between Rutherford's Atomic Theory and the Plum Pudding Model
   has to do with the fact that the Plum Pudding Model theorized that an atom was made up
   electrons (the plums) surrounded by a positively charged mass (the pudding). Rutherford
   later proved that this wasn't the case and theorized that atoms were comprised of a very
   small nucleus surrounded by electrons. The basic tenets of that statement are still held to
   be true in the modern era.

    Misconceptions

3. Many people believe that Rutherford's Atomic Theory is essentially the same as Bohr's
   Atomic Model, which is untrue. Bohr's Model was inspired by Rutherford's work and
   built upon it, but the Bohr Atomic Model contains certain elements that were definitely
   missing from Rutherford's Atomic Theory. The chief difference is the fact that Niels
   Bohr, when theorizing the design of the atom, included a thesis for the forces that help to
   hold the atomic structure together. The Bohr model is still widely accepted by
   contemporary scientists.

    Time Frame

4. The Bohr model of the atom was established in 1913 as the leading theory of atomic
   structure. Because of this, it can be truthfully said that Rutherford's Atomic Theory was
   an extremely short-lived phenomenon. Despite its brief time frame, however,
   Rutherford's theory was extremely important because Niels Bohr wouldn't have been able
   to develop his own model without the established background that Rutherford's Atomic
   Theory laid out for him.

    Effects
5. The effects of Rutherford's Atomic Theory are truly awe inspiring and the shockwaves
   are still being felt today. As mentioned, Rutherford's work paved the way for the Bohr
   Model of the atom, which is still widely accepted. Much of modern science and medicine
   has atomic theory at its very root. Without Rutherford's previous work, then, and without
   the later polishing work done by Niels Bohr, the face of contemporary science would be
   unimaginably different.




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