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History of Atoms
 Democritus was 1st to suggest existence of
atoms
 Said matter could be divided to a certain point
and then no more
 Dalton devised a theory based on Democritus’s
ideas in the early 1800s
 Theory had five points
History continued
 1. All matter is made of atoms
 2. Atoms can be distinguished by their mass
 3. Atoms of the same element have the same
mass and properties
 4. Atoms of different elements have different
masses and properties
 5. Atoms can’t be created, destroyed or
SUBDIVIDED
Atomic Structure
 Atoms are tiny units that determine the
properties of all matter.
 Atoms are the smallest unit of matter, though
they can be broken into smaller pieces.
 These pieces are the protons, neutrons, and
electrons.
Subatomic Particles
 Protons (p+) have a charge of +1.
They have a
mass of 1.673 * 10 –24 g.
 Electrons (e-) have a charge of –1. They have a
mass of 9.11 * 10-28 g.
 Neutrons (n) have no charge, but have the same
mass as a proton.
 Protons and neutrons are located in the nucleus
of an atom.
 Electrons orbit the nucleus, in energy levels.
Within an energy level, the electrons occupy
orbitals.
Electrons
 Discovered in 1897 by Thomson using CRT
 CRT = cathode ray tube
 Ran electricity through a gas, causing glowing
beam in tube
 Glow was due to motion of electrons among
energy levels
Protons
 Evidence for existence of protons seen in 1886,
Eugen Goldstein observed CRT with rays moving
toward CATHODE (- end) of tube
 Concluded must be +, since he knew cathode was –
and opposite charges attracted
 Was actually seeing movement of ions
 Existence of protons finally proven by Rutherford in
1919
Neutrons
 Non charged particles found in nucleus
 Existence proven in 1932 by James Chadwick
 He repeated experiments done by daughter &
son-in-law of Marie Curie, but with the goal of
finding a neutral particle
 He won Nobel Prize in 1935 for his discovery
Nucleus
 Discovered by Rutherford in “gold foil
experiment”
 Thomson had proposed p+ and e- were evenly
mixed throughout atom
 Rutherford shot α particles (He nuclei) at gold
foil, most passed through, but some
DEFLECTED
 Proposed model of atom containing mostly
empty space, with + charge and mass
concentrated in central region called NUCLEUS
Gold Foil Experiment
Rutherford expected the
α – particles (+ He
nuclei) to go through the
foil, or to rebound.
 Some particles were
deflected, apparently due
to repulsive force of like
charges

Niels Bohr (1913)
 Bohr suggested that e- were in orbits around the
nucleus and each e- has a certain energy that
determines its path around the nucleus.
 These paths are called energy levels. Electrons
can only be in certain energy levels.
 By 1925, Bohr’s model no longer explained all
observations.
 In Modern theory, it is suggested that e- behave
like waves on a vibrating string.
Atomic Math
 # of protons in an element = the element’s atomic
number.
 atomic # determines the element’s position in the
PT.
 All atoms are electrically neutral, so # of p+ must =
# of e Mass number = # of protons + # of neutrons.
 Isotopes result from differing #s of neutrons, thus
have different masses.
Atomic Mass
 Refer to relative masses of atoms since 1 atom
is EXTREMELY small
 Reference isotope is Carbon-12 (C atoms with 6
p+ & 6n)
 1 a.m.u. (amu) is 1 atomic mass unit, mass or 1
p+ or n
 Reported atomic masses are weighted
averages, based on the relative abundance of
natural isotopes
Example calculation
 Potassium has an atomic mass
of 39.14
 39K has a relative abundance of 93.3%, 40K has
a relative abundance of 0.012%, & 41K has a
relative abundance of 6.7%.
 (39*.933)+(40*.00012)+(41*.067) = 39.14
 (note #s may be off slightly from PT reported
value, due to rounding)