Atomic Theory
Historical Development [do not need to memorize names and dates]
NB: many discoveries were made independently by two scientists; attribution can be controversial
Year
Person(s)
Discovery
~400 BC
Democritus
Speculation on atomic nature of matter
1775
Lavoisier
Law of conservation of mass
1799
Proust
Law of definite proportions (law of constant composition)
~1800
Dalton
Law of multiple proportions
1807
Dalton
Proposed atomic theory of matter
- explained above “mass laws”
~1860
Mendeleev
Law of periodicity
Dalton theory
Modern theory
Relevant topics
Matter is made of indivisible
atoms.
Atoms are the smallest
particle of ordinary matter.
Subatomic particles
All atoms of a given element
are identical.
Elements exist as collections Isotopes and mass number
of isotopes
Atoms of different elements
have different relative
weights.
Mass v. weight
Isotopic abundances
Average atomic and molar masses
Atoms of different elements
may combine to form
compounds with simple
ratios.
Same
Laws of definite and multiple proportions
Empirical formulas
Atoms cannot be created nor True for physical and
destroyed.
chemical processes.
Law of conservation of mass
Nucleosynthesis
Atomic Theory
Discoveries of subatomic particles and related advances
During the 1800s, there was a gradual acceptance of a “billiard ball” model of an indivisible atom
Year
Person(s)
Discovery
1675-1835
Gas discharge, arc, and incandescent lamp development
1857-1895
Geissler tube, Crookes tube (1875), X-ray tube
~1886
Thompson & Goldstein
Proton discovered
1897
Thompson
Electron discovered
- same thing as cathode ray and beta ray
- leads to “plum pudding” model of atom
1899
Rutherford
Alpha (α) and beta (β) radiation discovered
1911
Rutherford
Proposed “nuclear model” of atom
1913
Moseley
Atomic number basis of periodicity not atomic mass
Bohr
Proposed “quantum model” of atom
De Broglie
Proposed dual wave-particle nature of matter
1927
Heisenberg
Proposed uncertainty principle
1932
Chadwick
Neutron discovered
The modern quantum model of the atom is quite abstract and math provides the best description. Between
wave-particle duality and the uncertainty principle, it is difficult to describe the nature of the atom without
some strange analogies. In this course, when discussing “locations” of electrons what is meant is the
probability density of the electron wavefunctions. Much more frequently we will discuss the energies of the
electrons. No more of that situation need concern us for now.
Atomic Theory
Below is a summary of what you should know about the structure of the atom.

Matter is composed of atoms.

Atoms are composed of protons, electrons, and neutrons.

Atom nucleus is composed of protons and neutrons.

Nucleus is surrounded by “cloud of electrons”.

Each element has a unique proton number.

Neutral atoms have equal proton and electron numbers.

Elements exist as collection of isotopes with varying neutron numbers.
Subatomic Particles
Z
atomic number
Z=P
A
mass number
A=Z+N=P+N
N=A-Z
Q
charge
Q=P-E
Q=0
Particle
Symbol Charge
(u)
(for neutral atom)
Charge
(C)
Mass
(u)
Mass
(g)
Mass
(kg)
Molar mass
(g)
Proton
p+
+1
+1.602×10-19
1.007
1.673×10-24
1.673×10-27
1.007
Neutron
n0
0
0
1.009
1.675×10-24
1.675×10-27
1.009
Electron
e-
-1
-1.602×10-19
0.000549
9.109×10-25
9.109×10-31
0.000549
Because the mass of the electron is ~1800 times smaller than the masses of the proton and neutron, the
electron mass is generally neglected for most calculations.
Thus, the mass of the atom is nearly the same as the mass of the nucleus. Also, because the masses of
protons and neutrons are ~1 atomic mass unit, the mass of a particular isotope is approximately equal to
the nucleon number.
Atomic Theory
Isotopes
Isotope
Symbol
P
N
A
Hydrogen-1
1
H
1
0
1
Hydrogen-2
2
H
1
1
2
Carbon-12
12
C
6
6
12
Carbon-13
Carbon-14
Average Atomic Mass
Weighted average of the isotope masses
Hydrogen
three isotopes exist
average mass = 1.008 g·mol -1
Carbon
eight isotopes exist
(only two are stable)
average mass = 12.01 g·mol -1
12
13
C
C
12.00
13.00
98.91 %
1.09 %