Transcript Atoms: The Building Blocks of Matter

Atoms: The
Building
Blocks of
Matter
Chapter 3
Foundations of the Atomic
Theory
If you crushed a sugar cube, you
would find a number of small
fragments that were still sugar.
If you crushed a fragment, you
would find a number of small
particles that were still sugar.
How long could you divide the
sugar until the particles were no
longer sugar?
This was the approach that
Democritus took in explaining
the atom. He said that you could
take a pair of shears and cut a
piece of copper in two and
sometime, you would reach a
piece that couldn’t be cut
anymore. He named this particle
an atom meaning indivisible.
In the late 1700’s, chemists had named a
number of elements. To them, an
element was a substance that could not
be broken down into simpler substances
by ordinary chemical means. With this
understanding, thousands of
experiments, and better technology,
scientists discovered several laws about
chemical reactions.
The Law of Conservation of
Mass states that matter is
neither created nor destroyed
in ordinary chemical or
physical reactions.
The Law of Definite
Proportions states that a
compound always contains
the same elements in the
same proportions by mass.
The Law of Multiple
Proportions states that when
two or more of the same
elements make up several
different compounds, the
ratios of the reacting
elements always compare in
the ratios of small whole
numbers.
Dalton’s Atomic Theory
In 1808, an English chemistry
teacher named John Dalton
proposed the atomic theory. From
data gathered in his student’s
experiments, he explained the
theories mentioned above and laid
the foundation for understanding the
atom.
Dalton’s Atomic Theory can be summed up in
the following statements:
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All matter is composed of small particles called
atoms.
Atoms of the same element are identical in size,
mass, and other properties; atoms of different
elements differ in size, mass, and properties.
Atoms cannot be subdivided, created, or
destroyed.
Atoms of different elements combine in small
number ratios to form compounds.
In chemical reactions, atoms are combined,
separated, or rearranged.
Modern Atomic Theory
Not all aspects of Dalton’s theory
have proven to be correct. We now
know that atoms can be subdivided
and that atoms of the same element
can have different masses. However
the most important concept that all
matter is made of elements and that
different elements have different
properties remains unchanged.
THE STRUCTURE OF THE
ATOM
As the atom was studied in depth, it
became clear that it contained subparticles. So today, we define the
atom as the smallest part of an
element that retains the chemical
properties of that element.
All atoms have two general regions,
a tiny, massive, positively charged
nucleus surrounded by a negatively
charged electron cloud. The
nucleus is made up of positively
charged protons and neutrons that
have no charge.
The Discovery of the Electron
Gases at atmospheric pressure do not conduct electricity
well, however, gases at very low pressures do conduct. A
glowing current will pass from a negatively charged cathode
to a positively charged anode. This stream was called a
cathode ray and the device was called a cathode ray tube.
Experiments using the tube showed that
an object placed in the tube cast a
shadow and the cathode ray could cause
a paddle wheel to roll along rails through
the tube. This indicated that the ray was
of a particle nature. Cathode rays are
deflected by a magnetic field and the rays
were deflected away from a negatively
charged object. This would indicated
that they carry a negative charge.
In 1897, Joseph John Thomson
measured the ratio of the charge
of the cathode particles to their
mass and found it always to be
the same. Thomson concluded
that cathode rays where made up
of identical, negatively charged
particles which were later named
electrons.
In 1909, Robert Millikan performed the
famous oil drop experiment to establish
the mass of an electron to be 1/2000 of a
hydrogen atom which was the smallest
particle known. Since, this has been
established at 9.109 x 10(-31) kg or 1/1837
of a hydrogen atom. By the same token,
Millikan determined the charge of an
electron.
Based on the discoveries made
about electrons, scientists made
two inferences:
• Because atoms are electrically
neutral, they must contain a positive
charge to balance the electrons.
• Atoms must contain other particles
that account for most of their mass.
Discovery of the Atomic Nucleus
In 1911, New Zealander Earnest Rutherford
and his associates Hans Geiger and Ernest
Marsden bombarded a thin gold foil with
alpha particles emitted from a radioactive
source. They expected an evenly charged
force field in the foil so they planned that the
particles would pass straight through. When
the detector was studied, they were greatly
surprised to find about 1 in 8000 particles
bounced straight back! Rutherford
exclaimed that this would be like firing a 15
inch artillery shell into a piece of tissue
paper and have it bounce back.
Rutherford concluded that
atoms must contain a very
small, dense, positively charged
nucleus. The nucleus of the
atom is very small. If it were the
size of a marble, the atom
would be larger than a football
field.
Composition of the Atomic
Nucleus
Except for the simplest type of hydrogen, the nucleus of atoms contains
two types of particles:
• They contain positively charged
protons. The number of protons
identifies the atom.
• They also contain neutrons which are
equal in mass to protons but do not
carry a charge.
Since most atoms in nature carry no net charge,
the number of protons and electrons in a resting
atom is the same. Otherwise, it would be
shocking!
Since particles with the same charge
repel each other, we would expect the
protons crowded together in the nucleus
to be unstable. However, when the
protons come into extremely close range
there is a strong attraction between
them. These strong, short range, protonproton, neutron-proton, and neutronneutron forces are referred to as nuclear
forces.
The radius of the atom is measured from
the center of the nucleus to the outer
edge of the electron cloud which
surrounds the nucleus. This distance
varies from 40 to 270 pm. Nuclei have
extremely high densities. In fact, the
density of a nucleus is about 200,000,000
metric tons per cubic centimeter.
Counting Atoms
Section Three
The atomic number of an element is the
number of protons in the nucleus. This
number identifies the element. It will be found
near the top of each grid on the periodic table.
For instance, hydrogen is 1, lithium is 3,
carbon is 6, and silver is 47.
The relative atomic mass of an element is
the mass of an atom of that element as
compared to an atom of carbon-12.
Since most elements have several
isotopes, this becomes a weighted
average (average atomic mass or atomic
mass) and may be followed by several
decimal places.
The mass number is the
number of protons and
neutrons in the nucleus. It is
derived by rounding off the
atomic mass.
Naturally occurring elements
may contain variations that have
differing numbers of neutrons.
Since they have the same
number of protons, they are still
the same element but they have
a different mass number. These
variations are called isotopes.
Most of the hydrogen in the world is
protium or hydrogen-1. This isotope of
hydrogen has one proton in the nucleus
surrounded by one electron. A small bit
of hydrogen, however, has a nucleus
containing one proton and one neutron.
This hydrogen-2 is called deuterium.
Some radioactive hydrogen or tritium
(hydrogen-3) has one proton and two
neutrons in the nucleus.
Relating Mass to Number of
Atoms
The mole is the SI unit for
measuring the amount of
substance. One mole is
precisely the amount of
substance that contains the
number of atoms in 12 grams
of Carbon-12.
This number of atoms is
called Avogadro’s Number
which is 6.022137 x 1023
atoms.
The molar mass of an element
is equal to the atomic mass
stated in grams.
It is important to understand
conversion mole quantities to
work efficiently in chemistry.