WAEC TUTORIAL:Atomic number, atomic mass, and isotopes...
Fundamental
properties of atoms including atomic number and atomic mass. The atomic number
is the number of protons in an atom, and isotopes have the same atomic number
but differ in the number of neutrons, thus they also have different mass number.
Introduction
Radioactivity
is actually a property of an atom. Radioactive atoms have unstable
nuclei, and
they will eventually release subatomic particles to become more stable, giving
off energy—radiation—in the process.
Often, elements come in both radioactive
and nonradioactive versions that differ in the number of neutrons they contain.
These different versions of elements are called isotopes, and small quantities
of radioactive isotopes often occur in nature. For instance, a small amount of
carbon exists in the atmosphere as radioactive carbon-14, and the amount of
carbon-14 found in fossils allows paleontologists to determine their age.
In this
article, we’ll look in more detail at the subatomic particles that different
atoms contain as well as what makes an isotope radioactive.
Atomic
number, atomic mass, and relative atomic mass
Atoms of each
element contain a characteristic number of protons. In fact, the number of
protons determines what atom we are looking at (e.g., all atoms with six
protons are carbon atoms); the number of protons in an atom is called the atomic
number. In contrast, the number of neutrons for a given element can vary.
Forms of the same atom that differ only in their number of neutrons are called isotopes.
Together, the number of protons and the number of neutrons determine an
element’s mass number: mass number = protons + neutrons. If you
want to calculate how many neutrons an atom has, you can simply subtract the
number of protons, or atomic number, from the mass number.
A property
closely related to an atom’s mass number is its atomic mass. The
atomic mass of a single atom is simply its total mass and is typically
expressed in atomic mass units or amu. By definition, an atom of carbon with
six neutrons, carbon-12, has an atomic mass of 12 amu. Other atoms don’t
generally have round-number atomic masses for reasons that are a little beyond
the scope of this article. In general, though, an atom's atomic mass will be
very close to its mass number, but will have some deviation in the decimal
places.
Since an
element’s isotopes have different atomic masses, scientists may also determine
the relative atomic mass—sometimes called the atomic weight—for
an element. The relative atomic mass is an average of the atomic masses of all
the different isotopes in a sample, with each isotope's contribution to the
average determined by how big a fraction of the sample it makes up. The
relative atomic masses given in periodic table entries—like the one for
hydrogen, below—are calculated for all the naturally occurring isotopes of each
element, weighted by the abundance of those isotopes on earth. Extraterrestrial
objects, like asteroids or meteors, might have very different isotope
abundances.
Isotopes
and radioactive decay
As mentioned
above, isotopes are different forms of an element that have the same number of
protons but different numbers of neutrons. Many elements—such as carbon,potassium, and uranium—have multiple naturally occurring isotopes. Carbon-12 contains six protons, six neutrons, and six electrons; therefore, it has a mass
number of 12 (six protons plus six neutrons). Carbon-14 contains six protons,
eight neutrons, and six electrons; its mass number is 14 (six protons plus
eight neutrons). These two alternate forms of carbon are isotopes.
Some isotopes
are stable, but others can emit, or kick out, subatomic particles to reach a
more stable, lower-energy, configuration. Such isotopes are called radioisotopes,
and the process in which they release particles and energy is known as decay.
Radioactive decay can cause a change in the number of protons in the nucleus;
when this happens, the identity of the atom changes (e.g., carbon-14 decaying
to nitrogen-14).
Radioactive decay
is a random but exponential process, and an isotope’s half-life is
the period over which half of the material will decay to a different,
relatively stable product. The ratio of the original isotope to its decay
product and to stable isotopes changes in a predictable way; this
predictability allows the relative abundance of the isotope to be used as a
clock that measures the time from the incorporation of the isotope (e.g., into
a fossil) to the present.
For example, carbon is normally present in the atmosphere in the form of
gases like carbon dioxide, and it exists in three isotopic forms: carbon-12 and
carbon-13, which are stable, and carbon-14, which is radioactive. These forms
of carbon are found in the atmosphere in relatively constant proportions, with
carbon-12 as the major form at about 99%, carbon-13 as a minor form at about
1%, and carbon-14 present only in tiny amounts^11start superscript, 1, end
superscript. As plants pull carbon dioxide from the air to make sugars, the
relative amount of carbon-14 in their tissues will be equal to the
concentration of carbon-14 in the atmosphere. As animals eat the plants, or eat
other animals that ate plants, the concentrations of carbon-14 in their bodies
will also match the atmospheric concentration. When an organism dies, it stops
taking in carbon-14, so the ratio of carbon-14 to carbon-12 in its remains,
such as fossilized bones, will decline as carbon-14 decays gradually to
nitrogen-14^22start superscript, 2, end superscript.
After a half-life of approximately 5,730 years, half of the carbon-14
that was initially present will have been converted to nitrogen-14. This
property can be used to date formerly living objects such as old bones or wood.
By comparing the ratio of carbon-14 to carbon-12 concentrations in an object to
the same ratio in the atmosphere, equivalent to the starting concentration for
the object, the fraction of the isotope that has not yet decayed can be
determined. On the basis of this fraction, the age of the material can be
calculated with accuracy if it is not much older than about 50,000 years. Other
elements have isotopes with different half lives, and can thus be used to
measure age on different timescales. For example, potassium-40 has a half-life
of 1.25 billion years, and uranium-235 has a half-life of about 700 million
years and has been used to measure the age of moon rocks^22start
superscript, 2, end superscript.
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