Name_______________________
I. The History of Atomic Theory
Atomic History, Structure, and Theory
Atomos (greek): Not to Be Cut
A. Greek philosopher Democritus began the search for a description of matter more than 2400 yrs ago.

His theory in short:
 Matter can’t be divided indefinitely, eventually the smallest indivisible piece would be obtained.
 He named the smallest piece of matter “atomos,” meaning “not to be cut.”
Philosophers of the time, Aristotle and Plato, had a more respected theory (4 elements = earth, fire, air, water).
B. Dalton’s Model
 In the early 1800s, the English Chemist John Dalton performed a number of experiments that eventually led to the
acceptance of the idea of atoms.
 Dalton’s Theory
 He deduced that all elements are composed of atoms.


Atoms are indivisible and indestructible particles.

Atoms of the same element are exactly alike.

Atoms of different elements are different.

Compounds are formed by the joining of atoms of two or more elements.
This theory became one of the foundations of modern chemistry.
C. Thompson’s Plum Pudding Model (Discovery of the electron)
 In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles.
 Thomson Model
 He proposed a model of the atom that is sometimes called the “plum pudding” model.
 Atoms were made from a positively charged substance with negatively charged electrons scattered about,
like raisins in a pudding. (or chocolate chip cookie model)
 Thomson studied the passage of an electric current through a gas.
 As the current passed through the gas, it gave off rays of negatively charged particles.
D. Rutherford’s Gold Foil Experiment (Discovery of the proton & nucleus)

(In 1908) Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold
foil (2000 atoms thick)
 Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil
without changing course at all.
 Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit
something solid. He knew that positive charges repel positive charges.
 This could only mean that the gold atoms in the sheet were mostly open space. Atoms were not a pudding
filled with a positively charged material.
 Rutherford concluded that an atom had a small, dense, positively charged center that repelled his
positively charged “bullets.”
 He called the center of the atom the “nucleus”
 The nucleus is tiny compared to the atom as a whole.
 Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The
negatively charged particles were scattered outside the nucleus around the atom’s edge.
 1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass
nearly equal to a proton
E. Bohr Model

In 1913, the Danish scientist Niels Bohr proposed an improvement. In his model, he placed each electron in a
specific energy level.



According to Bohr’s atomic model, electrons move in definite orbits around the nucleus, much like planets
circle the sun. These orbits, or energy levels, are located at certain distances from the nucleus.
The Wave Model

According to the theory of wave mechanics, electrons do not move about an atom in a definite path, like
the planets around the sun.

In fact, it is impossible to determine the exact location of an electron. The probable location of an electron
is based on how much energy the electron has.
According to the modern atomic model, an atom has a small positively charged nucleus surrounded by a large
region in which there are enough electrons to make an atom neutral.

Electron Cloud:
 A space in which electrons are likely to be found.

Electrons whirl about the nucleus billions of times
in one second

They are not moving around in random patterns.

Location of electrons depends upon how much
energy the electron has.

Depending on their energy they are locked into a
certain area in the cloud.

Electrons with the lowest energy are found
in the energy level closest to the nucleus

Electrons with the highest energy are found
in the outermost energy levels, farther from the nucleus.
Atoms and their Masses
A. Mass Number
B. Isotopes

mass # = protons + neutrons

always a whole number

not on the periodic table

Atoms of the same element
with different mass numbers.

Nuclear symbol

Hyphen notation = carbon-14
Mass #
Atomic #
Atomic # = # of protons in 1 atom of a given element
*In neutral atoms, the atomic # also equals the
number of electrons.
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Ex: Chlorine-37
Ex: F-19
atomic #:____17______
atomic #:____9______
mass #:_____37______
mass #:_____19______
# of protons:____17______
# of protons:____9______
# of electrons:____17______
# of electrons:____9______
# of neutrons:_____20______
# of neutrons:_____10______
C. Relative Atomic Mass

12C
atom = 1.992 ×
10-23

atomic mass unit (amu)
g
D. Average Atomic Mass

weighted average of all isotopes

on the Periodic Table

round to 2 decimal places
EX: Calculate the avg. atomic mass of oxygen
if its abundance in nature is 99.76% 16O,
0.04% 17O, and 0.20% 18O.
EX: Find chlorine’s average atomic mass if
approximately 8 of every 10 atoms are
chlorine-35 and 2 are chlorine-37.
1 amu
1p
1n
1 e-
= 1/12 the mass of a 12C atom
= 1.007276 amu
= 1.008665 amu
= 0.0005486 amu
Avg.
(mass)(%)  (mass)(%)
Atomic 
100
Mass
Symbol w/ charge:
Ion = an atom or molecule which has gained or lost one or more of its valence electrons , giving it a net
positive or negative electrical charge.
If an atom loses e-, the ion is pawsitive, a cation. This is because there are more p+ than e-.
If an atom gains e-, the ion is negative, an anion. This is because there are more e- than p+.
To find ion’s charge on an atom:
# p+ (positive charges)
+
# e- (negative charges)
=
ion’s charge
Ex: He+2
# of protons = +2
# of electrons = 0
Total charge
Symbol w/ Charge
= +2
# Protons
# Electrons
# Neutrons
Charge
Mass #
28
26
31
+2
59
17
17
19
0
36
(no charge)
7
10
7
-3
14