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Chapter 4
What is matter?
 Matter
is anything that has mass and
volume.
 Matter is made up of atoms.
 Energy is NOT matter and does not have
mass or volume.
 Energy and matter are related by E=mc2
Classifying Matter
H2O
Chemical vs. Physical Properties
Examples of Properties
 Carbon
reacts with oxygen to form carbon
dioxide and water.
 Water boils at 100ºC.
 Paper burns.
 A solution of KNO3 is colorless.
 Iron rusts.
 Solid sulfur is dull and yellow.
 Gold has a very high density (19.3 g/cm3).
Chemical vs. Physical Changes
Chemical vs. Physical Properties
Properties of matter are characteristics that can
be tested or observed and are used to identify
matter.
 Physical properties are those which can be
observed WITHOUT changing the chemical
make-up of the matter.
 Chemical properties can only be observed when
matter is involved in a chemical reaction, which
CHANGES the chemical composition.
 Chemical changes will create a new substance
whereas physical changes do not.
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Compounds and Diatomics
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Compounds are groups of atoms that are held
together by chemical bonds.
Compounds can be made by:
Ionic Bonding: Giving and taking electrons
(Metal-nonmetal bond)
Covalent Bonding: (Molecules) Sharing electrons
(nonmetal-nonmetal bond)
There are some gases that are bound to
themselves as elements (the diatomics):
Hydrogen = H2 Nitrogen = N2 Oxygen = O2
Fluorine = F2
Chlorine = Cl2 Bromine = Br2
Iodine = I2
How are compounds different
from mixtures?
A mixture is not like a compound in that the
parts of a mixture are NOT bonded together as
in a pure compound.
 Examples of compounds: water (H2O), sugar
(C6H12O6), carbon dioxide (CO2)…
 Examples of mixtures: air, steel, orange juice,
“Gorp” …
 Mixtures may be homogeneous (evenly
distributed in a single phase) or heterogeneous
(unevenly distributed in different phases).

Dalton 1803
Rutherford 1909
Bohr 1913
Thomson 1897
Present Day
Atomic History
The Greek philosopher Democritus
(400BC) coined the term atomon which
means “that which cannot be divided.”
 Idea of an indivisible thing that made
up all matter was refined by colorblind
chemist John Dalton in 1803. Among his

interests, Dalton was very interested in a
scientific explanation for his colorblindness
and the behavior of gases.
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Dalton published five principles of
matter:
Dalton’s Top Five
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All matter is made of indestructible and
indivisible atoms. (atoms are hard, unbreakable,
and the smallest thing there is)
Atoms of a given element have identical physical
and chemical properties. (all atoms of X will
These two are
behave the same anywhere)
usually combined
Different atoms have different properties. (X
behaves differently than Y)
Atoms combine in whole-number ratios to form
compounds. (two H’s and one O = Water (H2O)
Atoms cannot be divided, created or destroyed,
(just rearranged) in chemical reactions.
Daltons Laws:
Constant Composition: Ratios of atoms
in a compound is constant for that
compound. Ex: water is ALWAYS H2O
 Conservation of Mass: Mass is not
created or destroyed in a chemical
reaction.
 Multiple Proportions: Since atoms bond
in small, whole number ratios to form
compounds, their masses are small
whole number ratios. Ex: CO vs. CO2
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The CRT (cathode ray tube)
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A new invention, the cathode ray tube,
(c1850s) suggested the presence of positive and
negative charges.
This suggested that atoms must be divisible, and
Dalton’s theory had to be modified.
In 1897, English Physicist J. J. Thomson proposed
that the atom is a sphere of positive charge with
small areas of negative charge (he discovered
electrons).
This theory became known as the “plum pudding”
model after an English “dessert” of purple bread
and raisins.
Thomson’s cathode ray tube
and plum pudding model
electron beam
Millikan’s oil experiment
In 1897, Thompson used electrostatics
experiments to determine the charge-tomass ratio of 1.76 x 108 C/g.
 Millikan’s classic oil-drop experiment
allowed the charge of a single electron to
be determined: 1.60 x 10-19 C.
 Using these two numbers, we can calculate
.60 x10 C
 9.10 x10
the mass of an electron: m  11.76
x10 C / g

19
electron
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8
The mass of an electron is
about 1/2000 of the mass of a proton!
 28
g
Ernest Rutherford
While studying radioactive elements, New
Zealander Physicist Ernest Rutherford found
that radioactive alpha particles deflected
when fired at a very thin gold foil.
 This was known as the gold foil experiment,
and it suggested that the atom was not a
hard sphere as thought, but was mostly
space, with a small concentration of mass.
 This concentration of mass became known
as the nucleus.
 Link to experiment…
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Rutherford’s gold foil exp.
and new atomic model
The Bohr Model
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A Danish physicist, Niels Bohr, (a student of
Rutherford) rebuilt the model of the atom placing
the electrons in energy levels (the “solar system”
model).
Bohr was one of the founders of quantum physics
– a discipline that states that energy can be given
off in small packets or quanta of specific size.
Energy levels closer to the nucleus were lower in
energy than those farther away.
When a specific amount of energy was added to
an atom, an electron could jump into a higher
energy level.
Adding the Neutrons
Although theorized by Rutherford, British
physicist, James Chadwick, proved the
existence of massive, neutral particles.
 These particles came to be called
neutrons…
 and their discovery in 1932 opened the
door for more in depth investigations into
radioactive materials and gave the WWII
Allies the ability to enrich and purify
fissionable uranium, a necessity in the
production of nuclear weapons.
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The Modern Model
Democritus and
Dalton’s atom
Thompson’s electrons
electron
neutron
Rutherford’s space and
nucleus
proton
Bohr’s energy levels
(not to scale)
Chadwick’s neutrons
Most atoms ~ 1-5 Å
(Angstroms) = (1x10-10m)
What element is this? What is
is the mass of this isotope?
hey, does this
proton make my
MASS look big?
Elements
15.9994
We currently know of about 110
elements, 92 of which are naturally
occurring.
8
 We illustrate an element with an
OXYGEN
atomic symbol.
 The atomic number tells the number
atomic number
of protons and identifies the element.
 The atomic mass is the total
average atomic mass
mass of the protons plus the
neutrons. (on the P. T. the average
atomic mass is given)
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O
Weighted average mass
91.0%
25.2 kg
5.0%
16.1 kg
3.0%
10.0 kg
?%
12.0 kg
Natural Abundance - Isotopes
There is not just one type of each atom, there
are several. When an atom has more or less
neutrons than another atom of the same
element, we call them isotopes.
12C
13C
 For instance, the element carbon has 6
protons, but it could have 5, 6, 7, or 8
14C
neutrons, to form Carbon-11, Carbon-12,
Carbon-13, and Carbon-14. Each has a
different mass.
 In nature, there is a mix of different natural
isotopes. We use this mix to calculate average
atomic mass. ex: carbon is 12.011 amu
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Calculating Average Atomic Mass
To find average atomic mass, we multiply the
relative abundance of an isotope by the mass
of the isotope. We then add each of the
products for each isotope.
 Example: The isotopes of element Bob are
found below:
 Bob-18, 25%
0.25x18  0.60x19  0.15x20
 Bob-19, 60%
18.90amu
 Bob-20, 15%
 What is the average atomic mass of naturally
occurring Bob?
1amu = 1.66x10-24 grams
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Avg. atomic mass/mole atoms
Isotopes, Ions, and Allotropes
(Oh my)
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Isotopes are atoms of the same element with
different numbers of neutrons.
Ions are atoms of the same element with
different numbers of electrons. (Ions are easy to
create by adding or removing electrons from a
neutral atom).
Allotropes are forms of the same element, but
bonded in different structures.
Diamond and pencil graphite are examples of
allotropes. They are both pure carbon, but in
different structures.
Isotopes and allotropes
Ions
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An ion is an atom that has gained or
lost one or more electrons.
Recall that the octet rule predicts
that atoms try to achieve zero or
eight electrons in their outer
(valence) shell.
When an atom bonds with another
atom, it seeks to gain electrons or
lose them. For instance:
Cl has 7 and will gain one electron
Na has 1 and will lose one electron
Cl
Cl-
Na
Na+
Positive ions are
called “cations”
Negative ions are
called “anions”
Isotope symbols
+2
a charge is
shown here
for an ion
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Isotope symbols give the element symbol, the
mass and atomic numbers. From these, the
number of protons, neutrons, and electrons can
be determined.
Practice
How many protons, neutrons, and electrons
are present in (a) 27Al3+ (b) 79Se2 Write the isotope notation for an ion that
contains 20 protons, 21 neutrons, and 18
electrons.
 Write the isotope notation for an atom of
lead that has 128 neutrons.
 Write the isotope notation for the following:
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Why do nuclear reactions occur?
Nuclear reactions occur when a nucleus
becomes unstable.
 Protons and neutrons are attracted to each
other by the strong nuclear force. In a stable
nucleus, the attraction due to the strong force
is greater than the repulsion due to
electrostatic force. As elements get heavier,
they become more unstable. Extra neutrons
must be present to the nucleus (like glue) to
increase stability by increasing the strong force.
 Nuclear reactions are DIFFERENT from
chemical reactions, because new elements
form.
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What is radioactivity?
Radioactivity is the spontaneous emission of
radiation from an element to achieve a more
stable state.
 Uranium was the first radioactive element
isolated (by Bequerel), followed by radium
and polonium (by Marie Curie and her
husband Pierre).
 There are no stable isotopes for elements
after Bismuth (#83).
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Types of Radiation
Name
Symbols
Charge
Mass
Alpha
α or 42He
+2
4
Beta
β or 0-1e
-1
0
gamma
γ
0
0
Nuclear Reactions
Nuclear reactions change the composition of an
atom’s nucleus –the element will change!!
 Examples of naturally occurring nuclear
reactions include alpha and beta decay, and
fission and fusion.
 Some nuclei can become unstable by artificial
transmutation, where a nucleus is bombarded
(or shot) with a particle that creates instability
and causes radioactive decay.
 Nuclear reactions can produce enormous
amounts of energy as nuclear mass is
converted into energy (E=mc2)
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Radioactive Decay Equations
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Alpha decay equation.
parent isotope
daughter isotope
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Beta decay equation.
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Gamma decay involves energy
transitions (electromagnetic
waves), no particles are lost.
Other nuclear
reactions….fusion
and fission
2
3 H→ 4 He + 1 n
H
+
1
1
2
0
1 n→ 90 nuclear
143 fission
1 n
U
+
Sr
+
Xe
+
3
92
0
38
54
0
235
Nuclear Reactor Design
Nuclear Accidents
Hiroshima and
Nagasaki
Chernobyl
Three Mile Island
Products of Nuclear Reactions
In nuclear reactions, unstable nuclei change
their number of protons and neutrons.
 A DIFFERENT element is created by the
reaction, and large amounts of energy are
released (E=mc2), much more than in a
chemical reaction.
 Nuclear reactions result in the production of
new, more stable nuclei. Unstable nuclei will
continue to decay until a stable isotope is
produced.
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Radioactive Decay Series
A decay series ends with a stable isotope.
concentration
Radioactive decay and half-life
Nevada Nuclear Testing -1950’s
The Radium Girls
The Radium Girls were women
subjected to radiation exposure
at the United States Radium
Corporation factory, in Orange,
New Jersey around 1917. These
women painted the dials of watch
faces with luminous paint and
developed cancer (typically of the
mouth and jaw) as a result.
For fun, the Radium Girls painted their nails, teeth and faces with the deadly paint
produced at the factory, sometimes to surprise their boyfriends when the lights went out.
They mixed glue, water and radium powder, and then used camel hair brushes to apply
the glowing paint onto dial numbers. The going rate, for painting 250 dials a day, was
about a penny and a half per dial. The brushes would lose shape after a few strokes, so
the U.S. Radium supervisors encouraged their workers to point the brushes with their
lips, or use their tongues to keep them sharp.
THE END
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It is disconcerting to reflect on the
number of students we have flunked in
chemistry for not knowing what we
later found to be untrue.
– Quoted in Robert L. Weber, Science
With a Smile (1992)
Erwin Schrodinger
(1887-1961)
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Hi
Blame me! I started
all of this! In 1926 I
described the
location and energy
of electrons in an
atom with my
mathematical
model.