http://www.windows2universe.org/physical_science/physics/atom_particle/atomic_nucleus.html
Atoms are composed of a massive, central nucleus surrounded by a swarm of fast-moving electrons. The nucleus is made up of protons and, in most cases, neutrons. Almost all of the mass (more than 99%) of an atom is contained in the dense nucleus.
An atomic nucleus is much, much smaller than an atom. The cloud of electrons that "orbit" the nucleus and define the "size" of an atom is roughly 100,000 times as large as that atom's nucleus! For example, a helium atom has a size of about 1 Ångström (0.1 nanometers or 10 -10 meters), while its nucleus is only 1 femtometer
(10 -15 meters) in diameter. If you made a scale model of an atom with a nucleus the size of a pea, the electrons would zing around in a space larger than a major sports stadium! An atom is mostly empty space.
The number of protons in the nucleus determines what type of element the atom is. The number of protons is called the element's "atomic number". For example, hydrogen has an atomic number of one, since all hydrogen atoms have one proton in their nucleus. Carbon has 6 protons, so its atomic number is 6; oxygen has 8 protons, so its atomic number is 8. Uranium has 92 protons, so its atomic number is 92! If we count the number of protons plus neutrons, we get an atom's atomic mass. Most elements come in different versions, called "isotopes", with different numbers of neutrons.
For example, the most common form of carbon is carbon-12 (12C); that isotope of carbon has 6 protons and 6 neutrons, and thus an atomic mass of twelve.
Another isotope of carbon, carbon-14 (12C), has 6 protons and 8 neutrons, hence and atomic mass of fourteen. 12C is radioactive and is used to determine how old things are in a technique called "carbon dating".
Although the nucleus of an atom is far too small for us to see, here's one way of thinking about an atomic nucleus: as a cluster of tightly packed "balls". The cloud of electrons that "orbit" an atom's nucleus and define the "size" of an atom is roughly 100,000 times as large as that atom's nucleus!
Sometimes the electrons get stripped away from an atom. If an atom loses all of its electrons, leaving behind a "naked" atomic nucleus, the nucleus is called an ion. Ions moving at high speeds make up one type of particle radiation. These ions are usually made from relatively small nuclei, like the nucleus of a hydrogen atom (a single proton) or a helium atom (two neutrons and two protons). They can be much larger, though; some cosmic rays are very heavy ions from more massive atoms.
Introduction to Atoms
By Anne Marie Helmenstine, Ph.D., About.com Guide http://chemistry.about.com/od/atomicmolecularstructure/a/aa062804a.htm
All matter consists of particles called atoms. This is a list of the basic characteristics of atoms:
• Atoms cannot be divided using chemicals. They do consist of parts, which include protons, neutrons, and electrons, but an atom is a basic chemical building block of matter.
• Each electron has a negative electrical charge.
• Each proton has a positive electrical charge. The charge of a proton and an electron are equal in magnitude, yet opposite in sign. Electrons and protons are electrically attracted to each other.
• Each neutron is electrically neutral. In other words, neutrons do not have a charge and are not electrically attracted to either electrons or protons.
This is a (oversimplified) diagram of a helium atom, which has 2 protons, 2 neutrons, and
2 electrons.
• Protons and neutrons are about the same size as each other and are much larger than electrons. The mass of a proton is essentially the same as that of a neutron. The mass of a proton is 1840 times greater than the mass of an electron.
• The nucleus of an atom contains protons and neutrons. The nucleus carries a positive electrical charge.
• Electrons move around outside the nucleus.
• Almost all of the mass of an atom is in its nucleus; almost all of the volume of an atom is occupied by electrons.
Notes: Introduction to Atomic Structure 1 8/28/14

• The number of protons (also known as its atomic number ) determines the element. Varying the number of neutrons results in isotopes. Varying the number of electrons results in ions. Isotopes and ions of an atom with a constant number of protons are all variations of a single element.
• The particles within an atom are bound together by powerful forces. In general, electrons are easier to add or remove from an atom than a proton or neutron. Chemical reactions largely involve atoms or groups of atoms and the interactions between their electrons.
Since atomic number is the number of protons in an atom and atomic mass is the mass of protons, neutrons, and electrons in an atom, it seems intuitively obvious that increasing the number of protons would increase the atomic mass. However, if you look at the atomic masses on a periodic table you will see that cobalt (atomic number 27) is more massive than nickel (atomic number 28). Uranium (number 92) is more massive than neptunium (number 93). Different periodic tables even list different numbers for atomic masses. What's up with that, anyway?
The reason increasing atomic number doesn't always equate to increasing mass is because many atoms don't have a number of neutrons equal to the number of protons. In other words, several isotopes of an element may exist. If a sizeable portion of an element of lower atomic number exists in the form of heavy isotopes, then the mass of that element may (overall) be heavier than that of the next element. If there were no isotopes and all elements had a number of neutrons equal to the number of protons, then atomic mass would be approximately twice the atomic number (approximately because protons and neutrons don't have exactly the same mass... the mass of electrons is so small that it is negligible). Different periodic tables give differing atomic masses because the percentages of isotopes of an element may be considered changed from one publication to another.
Definition: An isotope is an atom with the same number of protons, but differing numbers of neutrons. Isotopes are different forms of a single element.
Examples: Carbon 12 and
Carbon 14 are both isotopes of carbon, one with 6 neutrons and one with 8 neutrons (both with
6 protons).
Parts of the Atom:
⦿
Outside
⦿
Whole Atom Properties
⦿
Center
• nucleus
• shells and orbitals
• electrons (negative charge)
• proton (positive charge)
• neutron (neutral)
• very small mass; 1/2000th of a proton
• Electrons DON’T follow specific orbits
• Electrons appear randomly in specific zones around the atom.
⦿
Atomic Number: number of protons
⦿
Mass Number: the total of neutrons and protons
⦿
Isotopes: Atoms with the same number of protons but a DIFFERENT number of neutrons.
⦿
Atomic Weight = weighted average of the masses of naturally occurring isotopes
⦿
Nuclear Notation – Symbols for writing nucleus information
⦿ P eriodic table information
• Mostly empty space
• Overall neutral charge
• Number of electrons = number of protons
1
2
1
3
Mass #
Atomi#
1
1
H
à Hydrogen
à
P+N
P
H
H
à Deuterium
à Tritium
Symbol
58
28
60
28
61
28
14
7
15
7
N
N
Ni
Ni
Ni
Atomic
Number
Unique to one kind of element
7
7
28
28
28
Mass Number Number of Protons
Varies depending on the number of neutrons
14
15
58
60
61
Unique to a specific element (exactly the same as atomic number)
7
7
28
28
28
Number of
Electrons
Equal to the number of protons for a neutral atom
7
7
28
28
28
Number of
Neutrons
Varies with different isotopes
7
8
30
32
33
% occurring in nature
Depends on the stability of an isotope.
99.63%
0.37%
68.27%
26.10%
1.13%
Notes: Introduction to Atomic Structure 2 8/28/14