Attempts at understanding matter date back to ancient Greek philosophers,
who viewed matter as being composed of elements, or simpler substances.
Two models were developed that considered matter to be
(1) continuous, or infinitely divisible, or (2) discontinuous, made up of
particles called atoms.
In the early 1800s, Dalton published an atomic theory, reasoning that
matter was composed of hard, indivisible atoms that were joined together
or dissociated during chemical change.
When a good air pump to provide a vacuum was invented in 1885,
cathode rays were observed to move from the negative terminal in an
evacuated glass tube.The nature of cathode rays was a mystery. The mystery
was solved in 1887 when Thomson discovered they were negatively
charged particles now known as electrons. Thomson had discovered the
first elementary particle of which atoms are made and measured their
charge-to-mass ratio.
Rutherford developed a solar system model based on experiments
with alpha particles scattered from a thin sheet of metal. This model had
a small, massive, and positively charged nucleus surrounded by moving
electrons. These electrons were calculated to be at a distance from the nucleus
of 100,000 times the radius of the nucleus, so the volume of an atom
is mostly empty space. Later, Rutherford proposed that the nucleus contained
two elementary particles: protons with a positive charge and neutrons
with no charge. The atomic number is the number of protons in an
atom. Atoms of elements with different numbers of neutrons are called
isotopes. The mass of each isotope is compared to the mass of carbon-12,
which is assigned a mass of exactly 12.00 atomic mass units. The mass
contribution of the isotopes of an element according to their abundance
is called the atomic weight of an element. Isotopes are identified by their
mass number, which is the sum of the number of protons and neutrons
in the nucleus. Isotopes are identified by their chemical symbol, with the
atomic number as a subscript and the mass number as a superscript.
Bohr developed a model of the hydrogen atom to explain the characteristic
line spectra emitted by hydrogen. His model specified that
(1) electrons can move only in allowed orbits, (2) electrons do not emit
radiant energy when they remain in an orbit, and (3) electrons move
from one allowed orbit to another when they gain or lose energy. When
an electron jumps from a higher orbit to a lower one, it gives up energy
in the form of a single photon. The energy of the photon corresponds to
the difference in energy between the two levels. The Bohr model worked
well for hydrogen but not for other atoms.
Schrödinger and others used the wave nature of the electron to develop
a new model of the atom called wave mechanics, or quantum mechanics.
This model was found to confirm exactly all the experimental data as well as predict new data. The quantum mechanical model describes the energy state of the electron in terms of quantum numbers based on the wave nature of the electron.The quantum numbers defined
the probability of the location of an electron in terms of fuzzy regions of space called orbitals.
The periodic table has horizontal rows of elements called periods and
vertical columns of elements called families. Members of a given family
have the same outer orbital electron configurations, and it is the electron
configuration that is mostly responsible for the chemical properties of an
element.
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