Atomic Structure

John Dalton

1808--Proposed the first atomic theory that all matter was composed of atoms and that only whole number of atoms can combine to form compounds. He also proposed that atoms of any one element differ in properties from atoms of another element.

J. J. Thomson

1897--Performed experiments in cathode ray tubes that proved the particles in the cathode ray were negatively charged. This led to the particles being named electrons.

Robert Millikan

1909--Determined the mass of an electron and confirmed that the electron carried a negative electric charge. Also proved that cathode rays have identical properties regardless of the element that produced them; therefore electrons are present in atoms of all elements.

Ernest Rutherford

1911--Because atoms are electrically neutral they must contain a positive charge to balance the negative electrons. Because electrons have so much less mass than atoms, atoms must contain other particles to account for their mass. Rutherford's Gold Foil experiment proved that positive charges occupied a very small section within the much larger area of an atom. He called this section the nucleus and determined that it was the site of the positive particles within the atom.

Neutron

a subatomic particle located in the nucleus of the atom. The neutron has a mass defined as 1amu and a charge of 0 (neutral)

Proton

a subatomic particle located in the nucleus of the atom. The proton has a charge of +1 (positive)

Law of conservation of mass

Matter is neither created nor destroyed during ordinary chemical reactions or physical changes

Law of definite proportions

The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.

Atomic mass unit

1 amu is exactly 1/12 the mass of a Carbon-12 atom. The mass number and relative atomic mass of a given nuclide are quite close, but not identical because the proton and neutron masses deviate slightly from 1 amu and also include the relatively small masses of all of the electrons. It is a unit created working with extremely tiny masses. 1 amu = 1.66 X 10-24 grams.

Electron

A negatively charged subatomic particle. Electrons exist outside of and surrounding the atom's nucleus. Each electron carries one unit of negative charge and has a very small mass as compared with that of a neutron or proton.

Nucleus

A small, heavy central portion of an atom, composed of protons and neutrons.

Valence electron

An electron in the outermost energy level that is available to be lost, gained or shared in the formation of chemical compounds.

Isotopes

Atoms of the same element that have different masses due to a varying number of neutrons. Isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons. In the example above, the atomic number for Mg is 12. The mass numbers are the top number and indicate the sum of the protons and neutrons. Although isotopes have different masses, they do not differ significantly in their chemical behavior.

Subatomic particle

Particles that are smaller than the atom

Atomic Number

The number of protons in the nucleus of each atom of that element. The atomic number identifies the element. Elements are arranged in the periodic table in order of increasing atomic number

Mass number

The total number of protons and neutrons in the nucleus of an isotope. Example: Carbon 14. Carbon has an atomic number of 6 = 6 protons. If the mass number is 14, carbon 14 must have 8 neutrons. Carbon 12 (mass number) would have 6 protons and 6 neutrons.

Average atomic mass

The weighted average of the mass of one atom of each naturally occurring isotope of an element in amus. The average atomic mass is weighted by the percentage at which each of an element's isotopes occurs in nature. The mass number and atomic mass of a given nuclide are quite close but not identical. The mass number is always a whole number, but the average atomic mass is always a decimal number for elements with stable isotopes. This is mainly due to the different number of neutrons in different isotopes but also due to including the tiny mass of the electrons and due to a tiny bit of mass being converted to binding energy (which holds the protons closely together even though they repel each other).

Atom

the smallest unit of an element that maintains the properties of that element. Nature's basic particle