A very simplistic visualization of an atom includes a dense nucleus comprised of a specific number of positively charged protons and uncharged neutrons surrounded by orbits of negatively charged electrons. One might be tempted to envision a configuration similar to a solar system with the nucleus representing the sun and the electrons representing planets. While this is a good starting point, you will soon discover that the solar system model does not provide a true representation of the atom.
The number of protons always equals the number of electrons in an atom, and that number is equal to the atomic number. For instance, carbon has an atomic number of six and therefore has six protons and six electrons.
The weight of an atom is determined by the number of neutrons and protons that are present in the nucleus. The proton and neutron, which are similar in mass, each weighs approximately 1,836 times greater than a single electron, thus the mass contributed by electrons is insignificant when determining atomic weight or atomic mass. The atomic mass is the sum of the protons and neutrons in the nucleus. Carbon has an atomic mass of twelve. Since there are six protons in carbon (remember, it has an atomic number of six and, therefore, must have six protons), it must have six neutrons:
Atomic Mass = # Protons + # Neutrons
The atomic mass of carbon = 12
The atomic # of carbon = 6 = the # of protons
# neutrons = Atomic Mass - # protons
# neutrons =12 - 6 = 6
While the number of protons and electrons remain constant in the neutral atom, the number of neutrons may vary within different atom species of the same element. As a result, the atomic mass for one atom may be different from another atom of the same element if the number of neutrons varies. Atomic mass must account for all possible species or nuclides (isotopes), of an atom. Carbon 12 with its 6 neutrons is by far the most common isotope of carbon. In reality, there is a carbon 14 which has eight neutrons and an atomic mass of 14. There is also a carbon 11 which has only five neutrons.
This is the sum of all nuclide (isotope), masses multiplied by their natural abundance. This weighted average is the relative mass listed in the Periodic Table. The relative percentage of each nuclide (isotope), appears to be similar throughout the world.
Since the number of protons (positive charges) always equals the number of electrons (negative charges) in an atom, positive charges equal negative charges and atoms in the elemental state have no charge. Only when an atom takes an electron from another atom does the particle become charged. This charged form of the atom is known as an ion. Positively charged ions are called cations, and negatively charged ions are called anions. For instance, when chlorine accepts an electron from sodium, the sodium ion that is formed will have one more proton than electrons. It will therefore have a positive charge and be called a cation. The chlorine (or chloride) ion will have one more electron than protons. It will take on a negative charge and be called an anion. The compound formed by this transfer of electrons is sodium chloride or table salt, which is nothing like the highly reactive sodium or extremely poisonous chlorine from which it was formed.
Until now we have focused on the nucleus. Lets turn our attention to the electrons, which surround the nucleus of the atom. Electrons are located in energy levels a term which has replaced the word shells, which was once used to describe the location of electrons. The word shell suggests a fixed position, which is far from reality. We will use energy level to describe the possible location of electrons.
There are seven energy levels. Each has a specific maximum number of electrons that can exist in it. The number of electrons, which an energy level can hold is equal to 2n2 where n = energy level. The letter n represents the principal quantum number that specifies the energy level of the atom in which an electron is located. The chart below identifies the various energy levels and maximum number of electrons possible. The energy level closest to the nucleus is represented by energy level 1.
|2n2||Possible # of electrons|
|5||2(52)||50 (theoretical, not filled)|
|6||2(62)||72 (theoretical, not filled)|
|7||2(72)||98 (theoretical, not filled)|
Within each principal energy level is one or more energy sublevels (Orbitals) or subshells. The number of sublevels possible for any one principal energy level is equal to the value of the quantum number (n) for that energy level. Note that I have purposely introduced the reader to the term quantum number quite early in this discussion. The idea is not to scare the reader, but rather to begin building up a certain comfort level with a term that may invoke a certain fear of the unreachable to the new chemist. While there are theoretically 7 possible sublevels, only four are actually used for the known elements. The others are not currently needed. Sublevels are numbered with consecutive whole numbers. The first sublevel is 0 followed by numbers 1 through 6. These numbers are the azimuthat quantum numbers, . The value of can never be greater than n-1. Based on the 112 known and verified elements, the following table represents possible sublevels in the atom (Jespersen, 1997).
|3||0,1,2||s, p, d|
|4||0,1,2,3||s, p, d, f|
|5||0,1,2,3||s, p, d, f|
|6||0,1,2,3||s, p, d|
|As more elements are identified, the sublevels of 5,6 and 7 will fill.|
If you need to cite this page, you can copy this text:
Roberta C. Barbalace. Anatomy of the Atom. EnvironmentalChemistry.com. Dec 1998. Accessed on-line: 9/19/2014