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Chapter 4
Notes sheet 1
What is matter?
 Matter
is defined as anything that has
mass and volume.
 Atoms are defined as the basic units of
matter.
NOTE: Energy is NOT matter and does not
have mass or volume. Energy and matter
are related by E=mc2
matter
pure substances
elements
mixtures
homogeneous
mixtures
compounds
one type of atom
different atoms
bonded to each other
uniform
(solution)
heterogeneous
mixtures
non-uniform
(separate parts)
Hg
I2
HgI2
element
element
compound
Classifying Matter
H2O
Compounds vs. mixtures



Compounds are groups of atoms that are held
together by chemical bonds. A molecule is the
smallest unit of a compound.
The elements within a compound can only be
separated by a chemical change, which breaks
the bonds in the compound.
Mixtures, however, can be separated by physical
methods such as filtration or evaporation, which
DO NOT chemically
alter the substances.
Compounds and Diatomics


Compounds are groups of two or more different
atoms that are held together by chemical bonds.
There are some gases that are bound to
themselves as pure elements (diatomic
molecules):
Hydrogen = H2
Nitrogen = N2 Oxygen = O2
Fluorine = F2
Chlorine = Cl2
Bromine = Br2
Iodine = I2
How would you classify these?
Chemical vs. Physical Properties
Chemical vs. Physical Changes
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. They are observed
during physical changes.
 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.

Chemical vs. Physical Changes
Chemical vs. Physical Changes
In a physical change, the chemical identity of
doesn’t change
the substance ___________________.
 In a chemical change, the chemical identity of
changes
a substance __________.
 Another name for a chemical change is a
______________
chemical _________________.
reaction

More 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).
Sheet 2
Matter is made
of atomos of air,
fire, wind and
earth!
460 BC
Rutherford 1909
Dalton 1803
Bohr 1913
Thomson 1897
Present Day
The Law of Conservation of Mass
Antoine Lavoisier (1700s) compared the masses
of substances before and after a reaction and
found that the mass always remained constant.
This led to the Law of Conservation of Mass.
In other words, matter cannot be created or
destroyed, only rearranged in chemical reactions.
The Law of Definite Proportions
French chemist
Joseph Proust
The same samples of a
pure compound always
contain elements in the
same mass proportion.
Dalton’s Atomic Theory
The idea of an
indivisible thing
that made up all
matter was
refined by
chemist John
Dalton in 1803.
His ideas were:
Daltons Law of Multiple Proportions:

Multiple Proportions: Since atoms bond
in small, whole number ratios to form
compounds, their masses are small
whole number ratios. Ex: CO vs. CO2
Thomson’s cathode ray tube
and plum pudding model
electron beam
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…

Millikan’s oil-drop experiment

Robert 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
the mass of an electron: The mass of an
electron is about 1/2000 of the mass of a
1.60 x10 C
proton!
m

 9.10 x10 g
19
electron
8
1.76 x10 C / g
 28
Rutherford’s gold foil exp.
and new atomic model
The Bohr Model
Neils Bohr, a Russian scientist, proposed
that the electrons must be in specific
energy “orbits” (called electron shells)
around the nucleus.
Lower
energy
shells
closer to
the
nucleus
The Modern Electron Cloud
(aka wave-mechanical) Model
Democritus and
Dalton’s atom
electron
Thompson’s electrons
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)
Sheet 3
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 mass number is the total
average atomic mass
number of the protons plus the
neutrons. (on the P. T. the average
atomic mass of isotopes is given)

O
Isotope symbols
a charge is
shown here for
an ion. This
+2
tells the
number of
electrons that
have been
gained or lost
 Isotope symbols give the element symbol, the
mass and atomic numbers. From these, the
number of protons, neutrons, and electrons can
be determined.
What element is this? What is
the mass of this isotope?
hey, does this
proton make my
MASS look big?
Ions




An ion is an atom that has gained or
lost one or more electrons.
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
Positive ions are
called “cations”
Cl
Cl-
Na
Na+
Negative ions are
called “anions”
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:

Formula mass (aka molar mass)
The formula mass for a compound is
atomic masses
the sum of the ______
______of
all elements in the formula.
Ex: Calculate the formula mass for
Al2(SO4)3
Notes 4
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

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.
Average atomic mass for one mole
of atoms
Isotopes, Ions, and Allotropes
(Oh my)




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
Sheet 5
Nuclear Chemistry and the
Band of Stability

#n > #p
What particles make up
the nucleus?
#n = #p protons and neutrons

What is the charge and
function of each
particle?
proton (+) -determines the
identity of the atom
neutron (0) –helps hold
nucleus together
Elements above
the band of
stability need to
decrease their
n/p ratio
Elements
below the
band of
stability need
to increase
their n/p ratio
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.

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).

Types of Radiation
Name
Charge
Mass
Alpha
+2
4
Beta
-1
0
0
0
+1
0
gamma
positron
Symbols
γ
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)

Radioactive Decay Equations
(natural transmutations)

Alpha decay equation:
parent isotope
daughter isotope

Beta decay equation (n→p+e-):

Gamma decay involves energy
transitions (electromagnetic
waves), no mass is lost!
Radioactive Decay Equations (natural
transmutations)

positron decay equation (n→p+e+):

Electron capture equation (n+e-→p):
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.

Radioactive Decay Series
A decay series ends with a stable isotope.
Sheet 6
How long does
radioactive material
last?
Half-life is defined as
the time it takes for
one half of the original
sample to decay.
 Some half-lives are
long, some are short…
(look at Table N)

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

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)

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.
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’s Top Five





All matter is made of indestructible and
indivisible atoms. (atoms are hard, unbreakable,
and the smallest thing there is)
Atoms of agiven 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.
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.

Bohr’s Atomic Model
Only certain orbits
are stable, and their
energies are related
to their radii.
An electron that absorbs
energy becomes
“excited” and jumps
from the “ground state”
to a higher energy orbit.
Light is emitted when the
electron returns to ground state.
Bohr’s Atomic Model

Atoms with electrons in the ground
state have 2, 8, 18, and 32 electrons
allowed in the first four shells.