The ATOM
A History of the Atom as brief as possible Democritus – matter
can be divided – coined ‘atom’
Aristotle – matter is continuous
Jump to 1700’s Mendeleev – organized
periodic table Dalton – measured pure
substances
Rutherford – discovered nucleus
Bohr – found pattern with atomic number
Schrodinger – movement of electrons
Planck and Heisenberg – nuclear chemistry
What the Heck? Atom comes from the Greek word
meaning indivisible.
Today we know that what we once knew as the atom, the smallest particle of life, is in fact divisible and contains smaller particles The Electron, Proton and Neutron
In an Atom: Protons and
Neutrons are in the Nucleus
Electrons are in the Electron cloud or energy levels.
In the Atom: Nucleus is positively
charged in the center of the atom
Electron cloud is negatively charged area surrounding the nucleus
To Illustrate
Dalton’s Laws
Still followed to this day: Law of conservation of mass: matter cannot
be created or destroyed in ordinary chemical or physical changes
Law of definite composition: a chemical compound contains the same elements in exactly the same proportions regardless of the size of the sample or the source
Modern Atomic Theory 1) All Matter is made up of very small particles
called atoms 2) Atoms of the same element have the same
chemical properties 3) While individual atoms of a given element
may not all have the same mass any sample of the element will have a definite average mass that is characteristic.
Modern Atomic Theory Cont.
4) Compounds are formed whenever two or more elements unite, with each atom loosing its characteristic properties as a result of the combination
5) Atoms are not subdivided in physical or chemical reactions
Protons, Neutrons and Electrons
Atomic Number = # of protons = # of Electrons Atomic weight = # of protons + # of neutrons
So if you had an atom of Lead
Atomic Number of lead = 82 Protons = 82 Electrons = 82
Atomic Mass of Silicon = 207.2 Neutrons = Mass-Protons = 207.2 - 82 =
125
Use Your Periodic TableName Symbol Atomic
MassProton Electron Neutron
Si
87.62
22
Iron
Bi
195.08
Zinc
77
71
35
Masses of Subatomic Particles Proton - 1.67265 x 10-24 g
Neutron - 1.67495 x 10-24 g
Electron - 9.10953 x 10-28 g
Electrons are 10,000 times smaller than Protons and Neutrons!
Isotope
Isotope - atoms with the same number of protons but different numbers of neutrons.
Number of protons make up the identity of the atom. I.E. anything with 6 protons is Carbon.
And why is the mass number a decimal? Carbon: (2 natural isotopes)
Carbon - 12 - • (total weight = 12)• (98.89 percent of Carbon atoms)
Carbon - 13• (total weight = 13)• (1.11 percent of Carbon atoms)
(12 x .9889)+(13 x .0111) = 12.01 the atomic mass of Carbon
Where the heck did that equation come from?
(12 x .9889) + (13 x .0111) = 12.01
(12 x .9889) the weight of Carbon - 12 times the percent of Carbon - 12
(13 x .0111) the weight of Carbon - 13 times the percent of Carbon - 13
Try it yourself:
Neon - 22 (Total weight = 22) (10% of Neon)
Neon - 20 (Total weight = 20) (90% of Neon)
Once More:
Hf - 176 = 5% Hf - 177 = 19% Hf - 178 = 27% Hf - 179 = 14% Hf - 180 = 35%