Chapter 3 - Atoms:
The Building Blocks of Matter
• There were two schools of thought of
the composition of the cosmos…
– is everything in the universe
continuous and infinitely divisible
– Or, is there a limit to how small you
can get?
• Particle theory was not the most
popular early opinion, but was
supported as early as Democritus
in ancient Greece.
From Philosophy to Science
• Democritus proposed that all the
matter is composed of tiny particles
called “Atomos”
– These “particles” were thought to be
indivisible
• Aristotle did not accept Democritus’
atom, he was of the “matter is
continuous” philosophy
– Because of Aristotle’s popularity
his theory was adopted as the
standard
From Philosophy to Science
• By the 1700’s nearly all chemists had
accepted the modern definition of an
element as a particle that is indivisible
• It was also understood at that time that
elements combine to form compounds
that are different in their properties than
the elements that composed them
– However, these understandings
were based on observations
not empirical evidence
From Philosophy to Science
• There was controversy as to whether
elements always combine in the same
proportion when forming a particular
compound.
– In the 1790’s, chemistry was
revolutionized by a new emphasis
on quantitative analysis because
of new and improved balances
• This new technology led
to the discovery of some new
scientific understandings
From Philosophy to Science
• The Law of Conservation of Mass:
– Proposed by Antoine Lavoisier
– States that mass is neither created nor
destroyed during ordinary chemical
reactions or physical changes.
– Which means the total mass of the
reactants must equal the total mass
of the products.
From Philosophy to Science

+
Carbon, C
Mass x
Oxygen, O
Mass y
Carbon Monoxide, CO
Mass x + Mass y

Carbon Monoxide, CO
Mass x + Mass y
+
Carbon, C
Mass x
Oxygen, O
Mass y
• The Law of Definite Proportions:
– The fact that 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
• NaCl is NaCl no matter if it is table
salt (small crystals) or rock salt
(large crystals)
From Philosophy to Science
• The Law of Multiple Proportions:
– If 2 or more different compounds
are composed of the same 2 elements,
then the ratio of the masses of the
2nd element combined with a certain
mass of the 1st element is always
a ratio of small whole numbers
From Philosophy to Science
+
Carbon
1
=
Oxygen
1
+
Carbon Monoxide,
1:1
=
Carbon
Oxygen
1
2
Carbon Dioxide,
1:2
• In 1808, John Dalton proposed an
explanation for each of the proposed
laws
– He reasoned that elements were
composed of atoms & that only whole
#’s of atoms can combine to form
compounds
– His ideas are now called the Atomic
Theory of Matter
Atomic Theory
1. All matter is composed of
extremely small particles called
atoms.
2. Atoms of a given element are
identical in size, mass, and
other properties; atoms of
different elements differ in size,
mass, & other properties.
ELEMENT
2
ELEMENT
3
ELEMENT
4
Atomic Theory
3. Atoms cannot be
subdivided, created, or
destroyed
4. atoms of different
elements combine in
simple whole # ratios to
form chem compds
5. in chemical rxns, atoms
are combined,
separated, or rearranged
Atomic Theory
+
+
• Through these statements, evidence
could be gathered to confirm or
discount its claims
– Not all of Dalton’s claims held up
to the scrutiny of experimentation
– Atoms CAN be divided into even
smaller particles
– Not every atom of an element
has an identical mass
Atomic Theory
• Dalton’s Atomic Theory of Matter has
been modified.
• What remains is…
1.All matter is composed of atoms
2.Atoms of any one element differ in
properties from atoms of another
element
• One of the disputed statements of
Dalton was that atoms are
indivisible
Atomic Theory
– In the 1800’s it was determined that
atoms are actually composed of
several basic types of smaller particles
– it’s the number and arrangement of
these particles that determine the
atom’s chemical properties.
• The def. of an atom that emerged
was, the smallest particle of an
element that retains the chemical
properties of that original
element.
Atomic Theory
• All atoms consist of 2 regions that
contain the subatomic particles
– The nucleus
– The electron cloud around the nucleus
• The nucleus is a very small region
located near the center of the atom
– In every atom the nucleus contains
at least 1 proton, which is positively
charged particle and usually
contains 1 or more neutral
particles called neutrons
Atomic Structure
• The electron cloud is the region that
surrounds the nucleus
– This region contains 1 or more electrons, which are negatively charged
subatomic particles
– The volume of the
electron cloud is much
larger than the nucleus
Atomic Structure
• The discovery of the first subatomic
particle took place in the late 1800’s.
– A power source was attached to two
metal ends of an evacuated glass tube,
called a cathode
ray tube.
– A beam of “light”
appears between
the two electrodes
called a cathode ray.
Discovery of the Electron
Electric Current
Electric Current
• Investigators began to study the ray
and they observed that…
1. An object placed in the path of the ray
cast a shadow on the glass
2. A paddle wheel placed in the path
of the cathode ray began to spin
3. Cathode rays were deflected
by a magnetic field
4. The rays were deflected away
from a negatively charged
object
Discovery of the Electron
• The first 2 observations support the
idea that the ray is composed of
tiny individual particles traveling
through the vacuum tube
• The second set of observations
support the evidence that the ray
is composed of a substance that
is negatively charged.
Discovery of the Electron
• J.J. Thomson studied the rays and proved
that they were tiny negative particles being
emitted from the metal atoms.
– Dubbed these tiny particles “electrons”
– And it was later determined that the
electrons were not part of the mass of the
atom.
Discovery of the Electron
• What can their work help us conclude
about the atom?
– atoms are composed of smaller
particles, and one of these components is negatively charged
– atoms are neutral, so there must
be an opposing (+) charge
– because electrons are essentially
mass-less, an opposing substance
makes up the mass of
the atom
First Atomic Model
• In 1886, E. Goldstein observed in the
cathode-ray tube a new set of rays
traveling in the opposite direction than
the cathode rays
– The new rays were called canal rays
and they proved to be positively
charged
– And the particles mass were about
2000 X’s that of the electron
Discovery of the Proton
• In 1932, the English physicist James
Chadwick discovered yet another
subatomic particle.
– the neutron is electrically neutral
– It’s mass is nearly equal to the proton
• Therefore the subatomic particles are
the electron, proton, and neutron.
Discovery of the Neutron
• Scientists still didn’t really understand
how the particles were put together in
an atom.
– This was a difficult question to resolve,
given how tiny atoms are.
• Most thought it likely that the atom
resembled Thomson’s model
Atomic Structure
Negative particles
embedded in a
sphere of positive
plasma-like matter.
THINK…
Chocolate Chip Cookie
• In 1911, Ernest Rutherford et al.
provided a more detailed picture
of the internal structure of the atom
• In his experiment, Rutherford directed
a narrow beam of alpha particles at a
very thin sheet of gold foil.
– Alpha particles (a) are He atoms
that have been stripped
of their electrons
Rutherford Model
• According to Thomson’s model,
the heavy, positive alpha particles
should pass easily through the gold,
with only a slight deflection
– And mostly that’s how it happened.
– However, they found 1 in every
8000 particles had actually been
deflected back toward the source.
Rutherford Model
• Rutherford suggested a new
structural model of the atom.
– He stated that all the positive charge
and the mass is concentrated in a
small core in the center of the atom,
AKA nucleus
– And that the atom is mostly empty
space with electrons surrounding
the positively charged nucleus like
planets around the sun.
Rutherford Model
Rutherford Model
• With the exception of Hydrogen, every
nucleus contains 2 kinds of particles
protons and neutrons
– they make up the mass of the atom (Mass
Number = Protons + Neutrons)
– Atoms are neutral because they contain
equal #’s of protons & electrons
Atomic Structure
• The atoms of different elements differ
in the # of protons in their nuclei and
therefore in their positive charge
– The # of protons the atom contains
determines the atom’s identity, also
known as atomic number.
• Only Oxygen contains 8 protons
• Only Fluorine contains 9 protons
Structure of the Atom
• The nucleus is composed of a densely
packed cluster of protons, which are all
electrically positive
– Don’t like charges repel?
– Why don’t they fly apart?
• When 2 protons are in very close
proximity, there is a strong force
of attraction between them.
– similar attraction exists
when neutrons are close
Structure of the Atom
• These short-range p+-n0, p+-p+, &
n0-n0 forces hold the nuclear particles
together, A.K.A strong nuclear forces.
– When these nuclear forces are strong
enough the atom is stable
– If the forces are not strong enough
the atom (heavier atoms) the atom
is unstable and becomes
radioactive.
Structure of the Atom
Ch 3.3: Atomic Number
• Elements are identified by the number of
PROTONS they contain
• The “atomic number” of an element is the number
of protons in the nucleus
– PROTONS IDENTIFIES AN ELEMENT!!!
• # protons in an atom = # electrons
– Because atoms are neutral!
Complete Symbol
Superscript →
Mass
number
Subscript →
Atomic
number
X
# OF PROTONS
+
# OF NEUTRONS
MASS
NUMBER
35
ATOMIC
NUMBER
17
Cl
NUMBER OF
PROTONS
Mass Number
Mass number is the number of
protons and neutrons in the nucleus
+ + n0
Mass
#
=
p
of an isotope:
p+
n0
e-
Mass #
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Element
Oxygen -
Mass number
18
Practice Problems
(1)Find the # of e-, p+ and n0 for sodium.
(mass # = 23)
Atomic # = 11 = # e- = # p+
# neutrons = 23-11 = 12
2) Find the # of e-, p+ and n0 for uranium.
(mass # = 238)
Atomic # = 92 = # e- = # p+
# neutrons = 238-92 = 146
Check for understanding:
n
If an element has 91 protons
and 140 neutrons find the:
a) Atomic number 91
b) Mass number 231
c) number of electrons 91
d) element name protactinium
Isotopes
• An isotope refers to atoms that have the same # of
protons, but a different number of neutrons.
• Because of this, they have different mass #’s.
Examples---> (1) Carbon-12 & Carbon-13
(2) Chlorine-35 & Chlorine-37
(Isotopes: The # after the name is the mass #.)
EXAMPLE OF AN ISOTOPE
ATOMIC MASS
Cl
Cl
35
37
17
17
18
NEUTRONS
20 NEUTRONS
ATOMIC NUMBER
Question #1
n
Find each of these:
a) Atomic number
b) Mass Number
c) number of protons
d) number of neutrons
e) number of electrons
80
35
Br
Question #2
n
If an element has an atomic
number of 34 and a mass
number of 78, what is the:
a) number of protons
b) number of neutrons
c) number of electrons
d) complete symbol
Atomic Mass
12
• Units = atomic mass unit (amu)
• The atomic masses listed in the Periodic Table are a
“weighted average” of all the isotopes of the
element.
Weighted Average
Practice Problems:
(1) In chemistry, chlorine has 2 isotopes:
Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance)
What is the weighted average atomic mass of chlorine?
35 x 0.758 = 26.53 amu
37 x 0.2423 = + 8.9651amu
35.4951 amu
Add them up!!!
(2) Oxygen has 3 isotopes:
O-16 (99.76%) O-17 (0.037%)
O-18 (0.2%)
Estimate oxygen’s average atomic mass.
Barely over 16.0 amu.
Relating Mass Numbers to Atoms
• The Mole: the amount of a substance that
contains as many particles as there are
atoms in exactly 12 grams of carbon-12.
• Avogadro’s Number: the number of
particles in exactly one mole of a pure
substance = 6.022 x 1023.
• Molar Mass: the mass of one mole of a
pure substance. Units = g/mol
• This is when we get to use dimensional
analysis!
• The conversion factors we need are:
1mol
23
6.022 x10 atoms
6.022x10 atom s
1m ol
and of course…molar mass
___ g
1mol
1m ol
___ g
23
Gram to Mole Conversions
___ g
1mol
Mass of
Element in
Grams
23
6.022x10 atom s
1m ol
Number of
Moles of
Element
1m ol
___ g
Number of
Atoms of
Element
1mol
6.022 x10 23 atoms
Practice Problem
• ALWAYS USE PARANTHESES AROUND
YOUR CONVERSION FACTORS!!
• You have 3.50 mol of Copper. What is it mass in
grams?