This section addresses, in whole or in part, the following Georgia GPS standard(s):
This section addresses, in whole or in part, the following Benchmarks for Scientific Literacy:
This section addresses, in whole or in part, the following National Science Education Standards:
All of the matter around you is made of atoms.
Atoms are the smallest particle of matter with constant properties.
All atoms are comprised of three types of subatomic particles:
All protons are exactly the same.
All neutrons are exactly the same.
All electrons are exactly the same.
Protons have a positive charge (+).
Electrons have a negative charge (-).
Neutrons do not have any charge. They are electrically neutral.
The nucleus of atom contains neutrons and protons.
Electrons orbit the nucleus.
Electrons are outside of the nucleus in electron shells that have different shapes at different distances from the nucleus.
In a stable atom, the number of electrons = the number of protons.
Mass of Subatomic Particles
Almost all the mass of an atom is concentrated in the nucleus.
Protons and neutrons have almost exactly the same mass.
The mass of a proton or neutron is 1.66 E -24 grams (1.66 x 10-24 g) or one atomic mass unit (AMU).
Electrons have a mass that is much less than that of the other particles (only about 1/1835 the mass of a proton). The mass of an electron is 9.05 E -28 grams (or 9.05 x 10-28 g). This number is a billionth of a billionth of a billionth of a gram.
Note: It is not possible for anyone or any machine that uses light to see a proton. The wavelength of light is too large to be able to detect anything as small as a proton.
The atom is mostly empty space.
Ernest Rutherford shot subatomic particles at a very thin piece of gold.
Most of the particles went straight through the gold. It was like shooting a rifle into a thin line of trees.
Some of the particles bounced off, some stuck inside, but the major portion of them passed through the gold foil.
By Rutherford's calculations, the nucleus in an atom is like a B-B in a boxcar.
There are only a few more than one hundred elements. Of those, only 83 are not naturally radioactive, and of those, only 50 or so are common.
Know the common elements well enough so that if you read or hear about one of them, you instantly know what it is.
Learn how to spell the names of the elements. Learn the symbols. Some of the symbols have one letter, some have two, but each element symbol has one and only one upper case letter in it.
You should know the name and symbol for the following elements. If you see the name, you should know the symbol. If you see the symbol, you should know the name. For some elements, there are other names, sometimes Latin, from which the element symbol was derived, which are listed in parentheses.
|Helium He||Lithium Li||Hydrogen H||Sodium (Natrium) Na|
|Boron B||Carbon C||Silicon Si||Calcium (Lime) Ca|
|Beryllium Be||Fluorine F||Neon Ne||Sulfur (Brimstone) S|
|Phosphorus P||Nitrogen N||Aluminum Al||Potassium (Kalium) K|
|Chlorine Cl||Argon Ar||Magnesium Mg||Iron (Ferrum) Fe|
|Bromine Br||Oxygen O||Manganese Mn||Copper (Cuprum) Cu|
|Cobalt Co||Nickel Ni||Chromium Cr||Lead (Plumbum) Pb|
|Zinc Zn||Krypton Kr||Rubidium Rb||Silver (Argentum) Ag|
|Iodine I||Platinum Pt||Cadmium Cd||Tin (Stannum) Sn|
|Cesium Cs||Barium Ba||Francium Fr||Antimony(Stibium) Sb|
|Bismuth Bi||Arsenic As||Strontium Sr||Tungsten(Wolfram)W|
|Radon Rn||Xenon Xe||Polonium Po||Gold (Aurum) Au|
|Radium Ra||Uranium U||Mercury (Hydrargyrum) Hg|
a. Each element is made up of tiny indivisible particles called atoms (* actually, atoms are divisible; they can be broken up into smaller particles -- some of these smaller particles are discussed below)
b. Atoms of a given element are identical; atoms of different elements are fundamentally different. (There may be some differences between atoms of the same element; we will discuss these differences below under isotopes.)
c. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms.
d. Chemical reactions involve the reorganization of the atoms, changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. (This is true except for nuclear reactions, but Dalton didn't know about nuclear reactions at the time.)
The atomic number is the number of protons in the nucleus of an atom. It is listed on the periodic table for each element. No two elements have the same atomic number (or the same number of protons), so the atomic number identifies the element.: (symbol: Z)
Mass number: (symbol: A) total number of protons and neutrons in the nucleus (not listed on the periodic table, since it varies)Atoms of the same element have the same atomic number, but may have different mass numbers.
Shorthand symbol notation for a particular atom (also called nuclide symbol notation):
AE (E = element's symbol; A = mass number; Z = atomic number)
represents a sodium atom which always has 11 protons and in this case has a mass number of 23.
(Note: This means that there are 12 neutrons. 23 - 11 = 12)
|Mass number - atomic number = # neutrons|
Write the symbol notation for the atom which has an atomic number of 9 and a mass number of 19.
How many protons and neutrons does this atom contain?
If this is a neutral atom, how many electrons does this atom contain?
What if this atom has a +2 charge -- how many electrons does it have?
Two atoms of the same element which have different numbers of neutrons are called different isotopes of that element. An isotope is an atom whose nucleus contain the same number of protons but a different number of neutrons.
The atomic mass (or atomic "weight") listed on the periodic table is the weighted average of the masses of all naturally occurring isotopes of an element in atomic mass units (amu).
Definition: 12 amu (exactly) = the mass of one atom of a 12C isotope; all other atomic masses are defined relative to this.
"Weighted average" means that the masses of each of the isotopes are "weighted" differently in calculating the average mass (just as quizzes might be weighted differently from tests or final exams in calculating your average grade). Masses of isotopes that are more abundant are weighted more heavily in the calculation of the average atomic mass for an element than masses of isotopes that are less abundant.
Example: There are four naturally occurring isotopes of iron. Below are listed the natural percent abundances and the masses of each of these isotopes:
|Isotope:||% Abundance:||Mass (amu):|
0.0582(53.94) + 0.9166(55.935) + 0.0219(56.935) + 0.0033(57.933) = 55.85 amu
The first isotope's mass of 53.94 amu is weighted as only 5.82% (0.0582) of the average mass, etc.
Don't worry about having to do these calculations; the results are already listed in the Periodic Table for you, but it's nice to understand where those values come from.
Content provided by Dr. Margaret Venable, Dr. Michael Denniston, and and Dr. David Wilner.
Some of this information is used with permission of Chemtutor.
Page created by Pamela J.W. Gore
Georgia Perimeter College,
Page created November 16, 2006