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– 400 B.C. thought matter
could not be divided indefinitely. First to
use the term ATOM (indivisible).
 ARISTOTLE- 350 B.C -modified an earlier
theory that matter was made of four
“elements”: earth, fire, water, air.
 his theory persisted for 2000 years.
Mass is neither created nor destroyed
during ordinary chemical reactions or
physical changes.
• This implies that for any chemical process in a
closed system, the mass of the reactants must
equal the mass of the products.
• Beginnings of the theory of conservation of mass
were stated by Epicurus(341-270 BC), who wrote
"the sum total of things was always such as it is
now, and such it will ever remain,".
 Joseph
Proust based on several
experiments conducted between 1798
and 1804 suggested the
Law of Definite Proportions- a chemical
compound contains the same elements in
exactly the same proportions by mass
regardless of the size of the sample or the
source of the compound.
 For example: oxygen makes up 8/9 of
the mass of any sample of pure water,
while hydrogen makes up the remaining
1/ of the mass.
 The
Law of Multiple Proportions- (John
Dalton) if two or more different
compounds are composed of the same
two elements, then the ratio of the masses
of the second element combined with a
certain mass of the first element is always
a ratio of small whole numbers.
 HUH?
 When
two elements can combine to form
more than one compound and the same
amount of the first element is used in each,
then the ratio of the amounts of the other
element will be a whole number.
 For example, carbon and oxygen combine
in carbon dioxide (CO2) and carbon
monoxide (CO). A sample of carbon dioxide
containing 1 gram of carbon contains 2.66
grams of oxygen; a sample of carbon
monoxide containing 1 gram of carbon
contains 1.33 grams of oxygen.
 The ratio of the two weights of oxygen
(2.66:1.33) is exactly 2:1
John Dalton 1808
All matter is made of tiny particles, called atoms.
Atoms are neither subdivided, created nor destroyed.
(Dalton based this hypothesis on the law of
conservation of mass)
Atoms of different elements combine in simple whole
number ratios, to form chemical compounds with
more than one ratio being possible for a given
combination of elements. (Dalton effectively
explained the law of definite proportions and law of
multiple proportions with this one)
Each element is made of a different kind of atom, and
the atoms of different elements have different masses.
In chemical reactions atoms are combined, separated
or rearranged. (Conservation of mass)
Atoms of a given element are identical in size, mass
and other properties. Atoms of different elements
differ in size, mass and other properties. (Oops!)
 Atoms
are divisible
 Isotopes exist
 Atomic theory has been modified but
these haven’t:
• All matter is composed of atoms.
• Atoms of any one element differ in mass and
properties from atoms of a different element.
 Read
Ch 3-1
 pg 69 q 1-3
 pg 87 q 2
 Chapter
The first evidence for
sub-atomic particles
came from experiments with
the conduction of electricity through gases
in sealed glass tubes at low pressures…
If an object is placed in the path of the cathode ray
then a shadow of the object is cast on the tube wall
at the end.
• Cathode rays travel in straight lines
The cathode ray can push a small paddle wheel up
an incline, against the force of gravity.
• cathode rays carry energy and do work.
The cathode ray is deflected from a straight line
path by a magnetic field.
• suggests that the two are related in some way.
J.J. Thomson succeeds in deflecting the cathode ray
with an electrical field.
• The cathode rays bend toward the positive pole,
confirming that cathode rays are negatively charged.
Thompson was able to measure the charge/mass ratio
of these electrons and found this ratio to be the same
regardless of what gas was in the tube or what metal
the electrodes were made from. (It was found that the
mass of the electron is extremely small when compared
to the charge)
His research proved the existence of electrons and
earned him the Nobel Prize for physics, 1906.
 1909
Oil drop experiment – showed that
the mass of an electron is about 1/2000
the mass of a hydrogen atom. Now
known to be 1/1837.
9.109 X 10-28g
 Confirmed that electrons carry a charge.
Millikan Oil Drop
 Nobel Prize in Physics, 1923
 Two
inferences about atoms could be
made based on what was learned about
the electron:
1. There must be a positive charge to balance the
negative charge of the electron.
2. Atoms must have other particles that account
for most of their mass
 1911
Ernest Rutherford, Hans Geiger, Ernest
firing of radioactive particles through
minutely thin metal foils (notably gold) and
detecting them using screens coated with
zinc sulfide (a scintillator).
 Rutherford found that although the vast
majority of particles passed straight through
the foil approximately 1 in 8000 were
deflected, leading him to his theory that
most of the atom was made up of 'empty
 This
model suggested that most of the
mass of the atom was contained in a small
bundle in the center of the atom and that
it had a positive charge.
 He called it the nucleus.
 When compared with the overall size of
the atom the nucleus is very small.
 Nobel
Prize in Chemistry 1908….but not
for the discovery of the nucleus.
 All
atomic nuclei are made of protons
and neutrons, except hydrogen.
OF THE ATOM, positive charge.
• Mass= 1.673 X 10-24g.
 Neutrons-
no charge, slightly greater
mass than a proton.
• Mass = 1.675 X 10-24g.
 Strong
Nuclear Force- a short-range
force responsible for binding of
protons and neutrons into atomic nuclei.
• One of the four fundamental forces of nature:
gravitational; electromagnetic; strong; weak
 Protons
and Neutrons are collectively
referred to as nucleons.
 Imagine
the nucleus to be the size of a
golf ball.
 On this scale the first electron shell
would be about one kilometer (0.62
miles) from the golf ball.
 The second shell- about four
kilometers (2.49 miles) from the golf
 The third -nine kilometers (5.59 miles).
 200 000 000 hydrogen atoms side by
side ≈ 1cm
 The
nucleus contains most (99.9%) of the
mass of the atom.
 Imagine a cube that is 1 mm on a side. If
filled with nuclear matter, it would have a
mass of about 200,000 tonnes.
• 1 tonne=2200 lbs
 some experiments have shown the nuclear force has a
"repulsive core," meaning that at very short distances,
the force switches from attractive to repulsive,
preventing the nucleus from collapsing in on itself.
The number of protons in the nucleus of the atom.
Shown above the symbol on the periodic table.
Ex: Silicon, Si, Atomic # 14
Find the atomic number of oxygen, O.
Silver, Ag.
 Atoms
of the same element that have
different masses.
 These
different masses are due to
different numbers of what atomic
 Examples
of Isotopes- Protium,
deuterium, and tritium are all
isotopes of hydrogen.
Number Number of
of protons neutrons
Number of
number of protons and
neutrons in the nucleus of an
 Relative
and Average Atomic
Because atoms are so small an arbitrary
value has been set as the standard.
Relative – AMU
A unit of mass equal to 1/12 the mass of
the most abundant isotope of carbon,
carbon- 12, which is assigned a mass
of 12 amu.
1.660540 x 10-24 g
Pg 85 q 2,3
Pg 87 q 19,20
mole of paper clips (3 cm long), chained
together, would wrap around the equator
450 trillion times.
 A mole of cantaloupes (6 inches in
diameter) has the same volume as Earth (1.1
× 1012 km3).
 A mole of pennies distributed evenly to
everyone in the world (7 billion people)
would enable each person to spend 1.6
million dollars per minute for an entire year.