ATOMIC THEORY
Building blocks of matter
Who are these men?
In this lesson, we’ll learn
about the men whose quests
for knowledge about the
fundamental nature of the
universe helped define our
views.
DEMOCRITUS
IN 400 BC, DEMOCRITUS
SAID:
 All matter is made of
tiny particles called
“atomos”
 Disputed by Aristotle
Why?
Eminent philosophers of
the time, Aristotle and
Plato, had a different
idea.
They favored the earth,
fire, air and water theory
of matter. They were
more popular, so the
atomos idea was buried
for approximately 2000
years.
Atomos
 To Democritus, atoms
were small, hard
particles like marbles
with different shapes
and sizes.
 Atoms were infinite in
number, always
moving and capable of
joining together.
For the next 2000 years…
 Alchemists tried to make gold from other
metals.
UNTIL…
1808 JOHN DALTON’S NEW ATOMIC
THEORY
LAWS FROM ATOMIC THEORY
(from last chapter)
 1. Law of conservation of matter—the mass of
the reactants before the reaction equals the
mass of the products after.
 2. Law of definite proportions—Every sample
of the same compound has the same mass
ratio of component elements.
 3. Law of multiple proportions—In a series of
compounds of the same two elements, the
ratio of an element in one compound to
another is also a small, whole number.
Dalton’s Model
 In the early 1800s,
English chemist
John Dalton
performed careful
experiments that
eventually led to the
acceptance of the
idea of atoms.
Dalton’s atomic theory
 1. All matter is made up of atoms
 2. Atoms of the same element are alike.
 3. Atoms of different elements are
different.
 4. Compounds have a definite
composition by weight and combine in
small whole number ratios.
 5. Atoms cannot be subdivided.
.
 This theory
became one
of the
foundations of
modern
chemistry.
Thomson’s Plum Pudding Model
 In 1897, the
English scientist
J.J. Thomson
provided the first
hint that an atom
is made of even
smaller particles.
Thomson Model
 His model of the
atom 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.
Thomson Model
 Thomson studied
passing an
electric current
through a gas.
 As the current
passed through, it
gave off rays of
negatively
charged particles.
 This surprised
Thomson,
because the
atoms of the gas
were uncharged.
Where had the
negative charges
come from?
Where did
they come
from?
He concluded that the negative
charges came from within the
atom.
A particle smaller than an atom had
to exist.
The atom was divisible!
Thomson called the negatively
charged “corpuscles,” today known
as electrons.
Since the gas was known to be
neutral, having no charge, he
reasoned that there must be
positively charged particles in the
atom.
But he could never find them.
Ernest Rutherford
 In 1908, the English
physicist Ernest
Rutherford was
hard at work on an
experiment that had
little to do with
unraveling the
mysteries of the
atomic structure.
 Rutherford’s experiment involved firing
a stream of tiny positively charged
particles at a thin sheet of gold foil
(2000 atoms thick)
Rutherford
– Most of the positive
particles passed through
the gold atoms in the foil
without changing course
at all.
– Some of the positive
charges did bounce
away from the gold sheet
as if they had hit
something solid. He
knew that like charges
repel.
Rutherford
 The gold atoms in the
sheet were mostly empty
space. Atoms were not a
plum pudding.
 Atom has a small, dense,
positively charged center
that repelled the positive
“bullets.”
 He called the center of the
atom the “nucleus”
 The nucleus is tiny
compared to the atom as a
whole.
http://chemmovies.unl.edu/ChemAnime/RUTH
ERFD/RUTHERFD.html
Atomic Particles
 Electron—discovered by Thomson in 1890’s
– Robert Millikan—determined the charge of an
electron in 1909 w/ oil drops
 Proton—discovered by Rutherford in 1911
 Neutron—discovered by James Chadwick in
1932
Particles and Charge
Mass
 Proton (p+) 1 amu
Charge
Location
+1
nucleus
0
nucleus
 Electron (e-) 1/1840 amu -1
electron
cloud
 Neutron (no) 1 amu
How Atoms Differ
Dalton said that all atoms of an element are
alike, but we know that is not completely
true. So what is alike?
1. All atoms of the same element have the
same number of protons.
2. If the atom is neutral, that means they
also have the same number of electrons.
3. The number of neutrons, however, can
vary.
How Atoms Differ
Atoms with the same atomic number are
the same element, but they may have
different numbers of neutrons.
 Atoms of the same element with a different
number of neutrons are called isotopes.
 Atoms of the same element with a different
number of electrons than protons are called
ions.
The Chemists’ Shorthand:
Atomic Symbols for Isotopes
#p+ + #no
Mass number 
Atomic number 
+
#p or
39
K
19
 Element Symbol
#e
Mass # - Atomic # = #no
Atomic Masses
 Elements occur in nature as mixtures
of isotopes
 Atomic mass is the weighted
average of all isotopes for an
element.
 Carbon =
98.89% 12C
1.11% 13C
<0.01% 14C
 Carbon atomic mass = 12.01 amu
MASS NUMBER AND AVERAGE
ATOMIC MASS
Atomic masses are based on CARBON.
The atomic mass unit is 1/12 of the mass
of one carbon atom.
How do we calculate average atomic
mass?
Multiply the % times the mass for each
isotope, then add them together.
Average atomic mass
 Calculate the average mass of isotopes
of neptunium with:
50.0% at 238.05 amu
29.4% at 235.1 amu
20.6% at 237.98 amu
(.500 x 238.05) + (.294 x 235.1) + (.206 x
237.98) = 237.17amu
Another problem:
Calculate the average atomic mass of
calcium with these isotopes:
28.4% at 40.06 amu
34.1% at 41.02 amu
22.8% at 40.89 amu
14.7% at 39.98 amu
(.284x40.06)+(.341x41.02)+(.228x40.89)+(.147x39.98)
40.56
One more for Arsenic
35.1% of 74.9 amu
18.6% of 74.2 amu
46.3% of 75.02 amu
74.83 amu
Atomic Mass
 Atomic mass is the
weighted average of
all of the known
isotopes of an
element, so will
always be shown as a
decimal number.
Covalent Chemical Bonding
 The forces that hold atoms together in
compounds. Covalent bonds result from
atoms sharing electrons between nonmetal
atoms.
 Molecule: a collection of covalently-bonded
atoms.
 Atom: representative particle for a
monatomic element
Ionic Chemical Bonding
Cation: A positive ion
Mg2+, NH4+
Anion: A negative ion
Cl, SO42
Ionic Bonding: Force of attraction between
oppositely charged ions. Smallest particle
called a formula unit.
Stupendous Seven
Periodic Table
Elements classified by:
properties
atomic number
Groups (vertical columns)—also called families
1A
2A
7A
8A
=
=
=
=
alkali metals
alkaline earth metals
halogens
noble gases
Periods (horizontal rows)
Periodic Table
Antoine Lavoisier , 1790’s made first list of known elements,
23 total. By 1870, there were 70!
John Newlands, 1864—Law of Octaves: When element were
placed in order of increasing atomic mass, every 8th
element repeated properties.
Lothar Meyer, 1869—Periodic table based on physical
characteristics only and increasing atomic mass.
Dmitri Mendeleev, 1869—Periodic table based on physical
and chemical characteristics and increasing atomic mass.
Predicted new elements.
Henry Moseley, 1913—Modern periodic law based on
subatomic particles: There is a periodic repetition of
chemical and physical properties of the elements when
they are arranged by increasing atomic number (protons).
Periodic Trends
Periodic Trends
Groups
Periods
Periodic Trends
Periodic Trends
Periodic Trends