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Transcript
Atoms : The Building Blocks of
Matter
Chapter 3
The Atom: From philosophical
idea to scientific theory

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
500 B. C. Aristotle: Matter was continuous
and there were no basic particles.
400 B. C. Democritus: Matter was composed
of particles, or atoms(Greek for indivisible)
1700’s: Elements could not be broken down
by ordinary chemical means. Elements
combined to form compounds that had
properties different from the original elements.
Such a transformation was called a chemical
reaction.
Foundations of Atomic Theory

1790’s Quantitative analysis came about and several
basic laws were discovered.

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Law of conservation of mass: Mass is neither created nor
destroyed in an ordinary chemical reaction or physical
change. (Lavoisier)
Law of definite proportions: A chemical compound contains
the same elements in exactly the same proportions by mass
regardless of the size of the sample or source of the
compound.
(Proust)
Law of multiple proportions: If two or more different
compounds are composed of the same two elements, then
the ratio of the masses of the second element combined with
a certain mass of the first element is always a small whole
number ratio. (Dalton)
Law of conservation of mass
Law of multiple proportions
Dalton’s Atomic Theory

1808 John Dalton proposed an explanation for the 3
laws. He stated the first atomic theory based on the
teachings of Democritus.

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All matter is composed of small particles called atoms.
Atoms of a given element are identical in size, mass, and
other properties; atoms of different elements differ in size,
mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.
Atoms of different elements combine in simple whole
number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or
rearranged.
Dalton’s Theory Explains the Laws

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The law of conservation of mass is explained by
the fact that chemical reactions involve
combination, separation, or rearrangement of
atoms and that during the process, no atoms are
lost, created, or subdivided.
The law of definite proportions results from the
fact that a given chemical compound is always
composed of the same combination of atoms.
As for the law of multiple proportions,
compounds of the same elements have the
same amount of one element and a multiple of
the second element.
Modern Atomic Theory

Not all aspects of Dalton’s atomic theory have
proved correct but Dalton is still known as the
father of atomic theory for his beginning
thoughts.

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We now know that atoms are composed of protons,
neutrons, and electrons. They do not have the same
properties as the atom however.
We also know that atoms of the same element exist
that differ in mass. Such atoms are called isotopes.
Atoms of different elements can have the same mass,
too.
The structure of the atom


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An atom is the smallest particle of an element
that retains the chemical properties of the
element.
All atoms consist of two regions---the tiny, dense
nucleus and the large electron cloud.
The nucleus contains protons(p+) and
neutrons(n0) while the electron cloud contains
electrons(e- ). These are subatomic particles.
The structure of the atom


http://ippex.pppl.gov/interactive/matter/ato
ms.html
http://www.pbs.org/wgbh/aso/tryit/atom/
Discovery of the electron


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Late 1800’s: The first subatomic particle, the
electron, was discovered in an evacuated
cathode ray tube by J. J. Thomson.
When an electric discharge was passed through
various gases at very low pressures in cathode
ray tubes, the surface of the tube opposite the
negative electrode(cathode) glowed.
Scientists hypothesized that the glow was due to
cathode rays which traveled from the cathode to
the positive electrode(anode).
Cathode Ray Tube
Cathode ray tube experiments

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An object placed between the cathode and
the opposite end of the tube cast a shadow
on the glass.
A paddle wheel placed on rails between the
electrodes rolled along the rails from the
cathode toward the anode.
Cathode rays were deflected by a magnetic
field.
Cathode rays were deflected away from a
negatively charged object.
Results of Cathode Ray Experiments

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All particles of the cathode ray are negative.
Thomson measured the charge to mass ratio.
He found that the ratio was always the same, no
matter what the cathode metal was or what the
gas was.
Thus Thomson concluded that all cathode rays
are composed of identical negatively charged
particles.
Millikan’s Oil Drop Experiment

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Millikan(1909) showed that the mass of the electron is
about 1/2000th the mass of a hydrogen atom or 9.11 x
10-31 kg.
He also found that the electron carries a charge of 1.6
x 10-19C.
Because all kinds of elements produced the same
cathode rays, he concluded that electrons were
fundamental to all atoms.
Because atoms are neutral, if there is a negative
charge, there must also be a positive one.
There must be other particles that help account for
mass of the atom.
Rutherford’s Gold Foil Experiment

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Rutherford(1911) bombarded thin gold foil
with alpha particles, which are like helium
nuclei.
Most the alpha particles passed through
the foil; others were deflected at angles.
Still others bounced back from the foil.
This experiment did more to characterize
the atom than any other.
Rutherford’s Experiment

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The atom was mostly space since most particles
passed through the foil.
The center of the atom was a small, very dense
core (nucleus) because there were so few
particles deflected.
The core was positively charged because it was
able to deflect the positive alpha particles.
He assumed that the electrons were outside the
nucleus orbiting the nucleus like planets around
the sun.
Rutherford’s Gold Foil Experiment
Discovery of the proton(1919)

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Protons were also discovered in a cathode ray
tube but the cathode in it was perforated.
Rays were found to come from the back of the
cathode. They were called canal rays.
They were found to be opposite in charge from
the cathode rays but equal in magnitude.
Rutherford and Goldstein concluded that a
positive ion may be formed from any neutral
atom so each atom must have positive particles.
Discovery of Neutrons

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There was extra mass in the atom that could not
be explained by protons and electrons.
In 1932, Chadwick found that there were neutral
particles in the nucleus that were given off as a
result of radioactive decay when Be atoms were
bombarded with alpha particles.
The mass of a neutron is approximately 1 g and
it may decompose into a proton and an electron.
The electron may be emitted as β rays while the
proton is reabsorbed by the nucleus.
Composition of the Nucleus

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All atomic nuclei are made up of protons and
neutrons except an isotope of hydrogen called
protium. It has no neutrons.
Protons are positive and electrons are negative
and equal numbers of them cause an object to
be neutral.
The number of protons determines an atom’s
identity.
Charges of the same type generally repel one
another but in the nucleus, protons are attracted
to one another and are held together with
gluons. These are short range nuclear forces.
Atomic Number

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The proton number is the atomic number,
Z. Elements on the periodic table are
ordered by atomic number.
The atomic number is an identifying
characteristic.
Moseley formulated the atomic number
concept.
Isotopes

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Isotopes are atoms of the same element
that have different numbers of neutrons.
The three isotopes of hydrogen are
protium, deuterium, and tritium.
Since there are different numbers of
neutrons, the masses of the isotopes will
be different.
Isotopes do not differ significantly in their
chemical behavior.
Mass Number

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Identifying an isotope requires knowing both the
name or atomic number and the mass of the
isotope.
The mass number, A, of the isotope is the sum
of the protons and the neutrons.


Specifying isotopes may be done in two ways.


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P+ + N0 = A
Write the name of the element, then a hyphen and
then the mass number. Carbon-13
Write the nuclear symbol. 235U is uranium-235.
Nuclide is the general term for any isotope of
any element.
Relative Atomic Masses

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Masses of atoms in grams are extremely small so
one atom has been arbitrarily chosen as the
standard and assigned a relative mass value. All
other masses are expressed in relation to this
standard.
The standard is the carbon-12 nuclide. It has been
assigned a mass of 12. Thus a hydrogen-1 atom
will have a mass of approximately 1/12th of carbon12 atom
One amu is exactly 1/12th the mass of a carbon-12
atom.
Average Atomic Masses of Elements

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Most elements occur naturally as mixtures
of isotopes.
The natural abundance of an isotope is
taken into consideration when determining
the average atomic mass of an element.
Average atomic mass is the weighted
average of the atomic masses of the
naturally occurring isotopes of an element.
Weighted Average
Do these
contribute
equally to
the sample
mass?
Relating Mass to Numbers of
Atoms

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The mole is the SI unit for amount of substance.
A mole is the amount of a substance that
contains as many particles as there are atoms in
exactly 12 g of carbon-12. It is a collective term
like dozen is.
The number of particles in a mole is 6.022 x
1023, which is Avogadro’s number.
Another definition of the mole is the amount of
substance in grams that contains Avogadro’s
number of particles.
The mass of one mole is the molar mass or
gram formula mass(gfm).