In our earlier discussion, we know that
everything in the universe is composed of atoms as proposed by Democritus
centuries ago. Further research showed that the atom is not as unbreakable as initially
thought of, but rather composed of still
smaller particles---protons and electrons. The early model of the atom consisted
of a positively charged protons inside a nucleus surrounded by a cloud of negatively
charged electrons. They reasoned correctly that for these atoms to continue to
exist, the positive nucleus must have the same amount of electrical charge as
the cloud of negative electrons orbiting around it, otherwise these atoms
disintegrated long time ago.
Different configurations of these
particles differentiate one element from another. They assigned these elements symbols
and atomic numbers. The atomic symbol is a one- or two- alphabetic characters and
the number corresponds to the number of protons which is also the same as the
number of electrons. Scientists started constructing the Periodic Table, a
2-dimensional array consisting of cubicles where to place the atomic symbols
relative to each other.
The smallest and the lightest element is
the hydrogen (symbol: H). It has an atomic number 1 because it has only one
proton and one electron orbiting around it. The electron is so tiny that its
mass is almost negligible and the mass of the atom is mostly due to the
nucleus. The heaviest naturally occurring element is uranium (symbol: U). It is
assigned atomic number 92 because it has 92 protons in its nucleus and 92
electrons around it.
The next lightest element is helium
(symbol: He) with 2 protons in its nucleus and 2 electrons orbiting around it.
In the periodic table, they assigned helium an atomic number 2. Because it has
2 protons, they assumed that helium atom is twice as heavy as hydrogen. But
when they compared the two elements, they found out that helium is 4 times
heavier than hydrogen! It means that helium has 4 proton-like particles inside
the nucleus but only two have positive charges to balance the two negatively
charged electrons! That is how neutrons were discovered. The neutron is a
particle as heavy as the proton but without electrical charge. The scientists
assign a new number to describe an element. They call it atomic mass unit
(amu). Hydrogen has an atomic number of 1 and an amu of 1. Helium has atomic
number 2 and an amu of 4 (2 protons plus two neutrons). The natural uranium has
an atomic number of 92 and an amu of 238 because it has 92 protons and 146
neutrons.
The universe is indeed a dynamic place.
Billions of stars are continuously burning themselves, their nuclear furnaces
spewing out charged and uncharged particles and rays of energies in all
directions. In this pretty busy place, an atom or a group of atoms may catch
extra neutrons or lose some. So it is
not unusual that atoms of the same element may not have the same neutrons in
their nuclei. Atoms of the same element but don’t have the same number of neutrons
are called isotopes. Every element found in the periodic table has isotopes.
Isotopes of an element are described by the atomic symbol and the amu which is
slightly higher or lower than the amu of the natural element. For example, natural
carbon atom has 6 protons and 6 neutrons and is usually referred to as C-12.
But some carbon atoms hold 2 more neutrons and it is referred to as C-14. Another
good example is hydrogen. Some hydrogen atoms contain an exra neutron and are
designated 2H while some contain 2 neutrons and are designated 3H. Hydrogen
isotopes are given special names. Thus 2H is called deuterium and 3H is called
tritium.
Most isotopes are unstable. As such, they
try to eject excess particles and radiate energies until they return to a
stable state. The process of ejecting
those excess particles and energies is called radioactivity.
The transformation from being an unstable isotope
to a stable substance is called a decay.
The time it takes for a radioactive element to decay to half of its original
amount is called half-life. The half life of a radioactive isotope is constant
regardless of what the original amount was and is a distinct characteristic of
that particular isotope. For example, Iodine 131 has a half-life of 8 days. If
at the start you have 2 grams, 8 days later, what is left is 1 gram, another 8
days, one-half gram is left, another 8 days later, one-fourth gram is left, and
so on… Radium 226, on the other hand has a half-life of 1,600 years!
Radioisotopes have many uses in our modern
world. The half-life of an isotope in a substance can be used to measure long
periods of time that help our scientists determine events that happened
thousands or even millions of years ago.
