1. The Atom

2. Dalton’s Atomic Theory
1. All elements are composed of tiny indivisible particles called atoms 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 3. Atoms can physically mix together or can chemically combine in simple whole number ratios. 4. Chemical reactions occur when atoms are separated , joined or rearranged. Atoms of one element never change into atoms of another element as a result of a chemical reaction.

3. Electrons Electrons are negatively charged particles in an atom
JJ Thomson passed current through gases at low pressure and found cathode rays were attracted to the metal plates which had a positive charge. Plates with a negative charge repelled the rays, so a cathode ray is a stream of negatively charged particles called electrons.
Millikan found the quantity of charge carried by an electron and the ratio of charge to mass determining the mass of an electron.

4. Protons and Neutrons
Goldstein discovered the proton using the cathode ray, a particle with a positive charge. The proton mass is calculated as 1840 times the mass of an electron.
Chadwick discovered a neutral particle called the neutron with the same mass as a proton.
Rutherford directed a narrow beam of alpha particle at a thin sheet of gold foil. Most of the alpha particles went through the foil with only a few being repelled because an atom has a very small positive nucleus. He concluded an atom is made up of mostly empty space.

5. Atomic Number
Protons and neutrons are in the middle of the atom and have the most mass. Electrons surround the nucleus with little mass.
The atomic number is the proton number and electron number in a neutral atom.
Neutrons must always be calculated. Neutrons + protons =Mass
Mass – protons = neutrons
On your periodic table Atomic weight = Mass (Average atomic mass, atomic mass unit)

6. Mass Mass may be represented as……
Hydrogen-1 or H-1, Gold-197 or Au-197, Carbon-12 or C-12
Or 1H, 197Au or 12C
Proton numbers can be represented as…..
0H, 79AU, 6C
Together superscript to the left is mass, subscript to the left is atomic/proton number 10H, Au, 126C

7. Isotopes
Protons remain the same, neutrons differ. H-1, H-2, H-3 is an example
Average atomic mass/atomic mass is based on isotopes and is the weighted average mass of an atom in a naturally occurring sample of the element.
Atomic mass unit is defined as 1/12 mass of the carbon-12 atom, protons + neutrons = mass
Average atomic mass considers abundance

8. How to calculate Average Atomic Mass
Mass X percent abundance = percent abundance mass
Add all of the percent abundance mass’ = overall percent abundance mass
Overall percent abundance mass / 100 = average atomic mass
X = multiply
/ = divide