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The Atom
Atomic Models
• Think with me…about sugar crystals,
you can see that they are small
crystals and every crystal is identical.
– You may grind these particles into a
very fine powder, but each tiny piece
is still sugar.
• If you dissolve the sugar in water,
the sugar particles become
virtually invisible.
From Philosophy to Science
– You could even look at the sugar
solution under a microscope and
you’d still not be able to see
the sugar.
• However, you know it is still there
because you can taste it.
• These kind of observations and logic
patterns led ancient philosophers
to ponder the design
of the universe
From Philosophy to Science
• 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 rxns 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
Copper Carbonate
• 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
Steering Coils
The electron beam is painting all 525 lines 30 times
per sec, it paints a total of 15,750 lines per sec.
• 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 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
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”
• Robert Millikan then used an ingenious
investigation to calculate the mass
to charge ratio of an atom
– He 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 E’s are essentially
mass-less, an opposing substance
that makes up the mass
of the atom
First Atomic Model
Negative particles
embedded in a
sphere of positive
plasma-like matter.
THINK…
Chocolate Chip Cookie
Cathode Ray Tube
• 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
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
The Neutron
electron
e-
-1
0
9.11x10-28
proton
p+
+1
1
1.67x10-24
neutron
n0
0
1
1.67x10-24
Structure of the Atom
• 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
• 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
Rutherford Model
• Rutherford’s planet system model was
an improvement over earlier models,
but it was still not complete.
– Physics says that electrons can’t orbit
the nucleus without losing energy,
• Losing energy would cause the
electron to spiral into the nucleus.
• The attraction of the electron to the
nucleus would cause it to spiral into
the nucleus as well
Rutherford Model
• Niels Bohr proposed a new model
that would allow the electrons to be
outside the nucleus and in orbit
around the nucleus.
– His model coupled Rutherford’s model
with a new concept of energy in
Physics called quantum mechanics,
• Bohr proposed that the electrons
aren’t on any random orbit around
the nucleus, they are on
“special” orbits
Bohr Model
• Bohr’s Model restricts the orbits on
which an electron can be
– The bases for what orbit an electron is
allowed is entirely based on how much
energy the electron has
• If it has any more energy or any less
energy it would be forced to be on a
different path of different energy
– The energy of the electron is
quantized, which means it is
of a very specific quantity
Bohr Model
– Each path or level of energy that the
electron is on is given a label of “n”
• Such that n=1 is the closest energy
level to the nucleus
• n=2 is higher in energy and outside of,
but adjacent to n=1, and so on…
–Each energy level can only hold
a certain number of electrons (2n2)
• n=1 can hold 2 electrons
• n=2 can only hold 8 electrons
• n=3 can hold 18 electrons
• Etc.
Bohr Model
Bohr Model of the Atom
Model describes the paths
of electrons as energy levels.
The electrons are only allowed
to have a certain amount of energy
which restricts their path around
the nucleus.
Bohr Model
Bohr 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)
• Proton has a charge equal to but
opposite of the charge of an elec.
– 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
• Only Oxygen contains 8 protons
• Only Fluorine contains 9 protons
• Only Neon contains 10 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
Nuclear Forces
The Quark
• Basic Truth: All atoms contain the
same basic parts, but atoms of
different elements have different
numbers of protons.
– The PT lists atoms in consecutive
order by their Atomic Number (Z)
– The atomic number is directly
related to the number of protons
in the nucleus of each atom
that element
Counting Atoms
of
• The atomic number is found above
the elemental symbol on the PT and
it defines the type of element
– Atomic #47 can only be Ag and it also
can only have 47 protons in each
nucleus
– Because atoms are neutral,
we know from the atomic number
the atom must also contain
47 electrons.
Counting Atoms
• The total number of protons &
neutrons determines the mass of the
atom
– Called the Mass Number
– A Carbon atom, has 6 protons and 6
neutrons, so its mass number is 12
• If you know the atomic number &
mass number of an atom of any
element, you can determine the
atom’s composition
Counting Atoms
ATOMS OF THE 1ST TEN ATOMS
NAME
SYMBOL
ATOMIC #
p+
n0
MASS #
e-
Hydrogen
H
He
Li
Be
B
C
N
O
F
Ne
1
2
3
4
5
6
7
8
9
10
1
2
3
4
5
6
7
8
9
10
0
2
4
5
6
6
7
8
10
10
1
4
7
9
11
12
14
16
19
20
1
2
3
4
5
6
7
8
9
10
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
• Every Cl atom has 17 protons, w/o
exception, but not every Cl atom has
18 neutrons.
– Atoms with the same # of protons but
contain different #s of neutrons are
called isotopes.
• Since isotopes of an element have
different #s of neutrons they have
different masses
Counting Atoms
• Isotopes are chemically alike because
they have identical numbers of protons
and electrons
– It’s the electrons and protons that are
responsible for chemical behavior
• Isotopes can be noted using hyphen
notation (Cl-35 vs Cl-37)
– elemental symbol hyphen mass
number
Counting Atoms
or Na-23
or Na-24
Isotope: one of two or more atoms having
the same number of protons but
different numbers of neutrons
Name
Symb
Atomic Mass
#
#
# p+
# e-
# n0
Isotopic
Symbol
Strontium
W
82
70
74
33
Practice
As