Chapter 3: Organization of Matter

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Chapter 3 Organization of Matter

• The Greek thinkers, Aristotle and Democritus
  proposed two conflicting theories of matter.
• Democritus Matter is made of the small
  indivisible particles known as atoms.
• Aristotle Matter is continuous and uniform
  throughout and consists of four elements earth,
  air, fire, and water.
• Neither theory was based on empirical
  (experimental) facts or data.

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John Dalton 1803

• John Dalton was an English schoolteacher and
  developed an understanding of the following laws
  
• Law of conservation of mass.
• Law of definite composition (proportions)
• Law of multiple proportions.
• Dalton used Democrituss idea of the atom and
  experimented to prove the validity of his five
  proposals.

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Daltons Laws 1

• Atoms and the conservation of matter.
• Mass of all atoms involved in a chemical reaction
  is always conserved.
• Since chemical reactions involve the combing,
  separating and rearranging of atoms, then matter
  cannot be created or destroyed.
• Atoms and the definite composition of chemical
  compounds.
• Atoms that combine to form a new chemical
  compound have proportionality by mass.
• Water. H20, 1 gram of Hydrogen reacts with 8
  grams of Oxygen.

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Daltons Laws 2

• Law of Multiple Proportions
• Dalton proposed that some elements could
  combine to form two or more different chemical
  compounds.
• Examples are
• 1. Carbon and Oxygen
• CO -carbon monoxide
• CO2 -carbon dioxide
• 2. Hydrogen and Oxygen
• H2O -water
• H2O2-hydrogen peroxide

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Daltons Laws 3

• Law of Multiple Proportions
• If two or more different compounds are made of
  the same two elements, the masses of these
  elements can be expressed as ratios of small
  whole numbers.
• John Dalton began the inquiry on the nature of
  matter and the use of experimentation to prove or
  disprove theories.

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Daltons Proposals

• Five major proposals of Dalton's theory.
• All matter is made up of extremely small
  particles called atoms.
• Atoms of the same type of matter are identical in
  size, mass and other properties.
• Atoms can not be broken down into smaller
  particles, or be created or destroyed.
• Atoms of unlike elements combine in simple whole
  number ratios to form chemical compounds.
• In chemical reactions, atoms are combined,
  rearranged or separated.

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Atomic Structure

• All atoms are made of three fundamental subatomic
  particles
• Electrons-subatomic particles that orbit around
  the central core or nucleus.
• Protons- subatomic particles found in the nucleus
  of an atom.
• Neutrons- subatomic particles found in the
  nucleus of an atom.

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Description of Electrons

• 1. Negatively charged subatomic particle.
• 2. Very little mass, 9.1 x 10 -28 grams.
• 3. Large charge to mass ratio. (compared to
  other subatomic particles.)
• 4. Electrons of one element are identical in
  composition to electrons in another element.
• 5. Electrons orbit in all directions around, the
  nucleus of an atom.
• 6. The mass of one electron is approximately
  1/1840 of the mass of a proton.

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Experimental evidence of the electron
Movement of the paddlewheel gave evidence that
electrons must have mass.
1897 J.J. Thomson cathode rays are negatively
charged particles.
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Robert Millikans Experiment
Millikans experiment measured the mass of an
electron. Was conducted in 1909, and measured the
mass of an electron to the power of ten, 10-28
grams. Provided the evidence that would be used
by others to propose atomic theory.
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Description of the nucleus

• 1. Particles found in the nucleus are referred
  to as NUCLEONS.
• 2. Two main types of nucleons are Protons and
  Neutrons.
• 3. Overall charge of the nucleus is positive.
• 4. Densely concentrated mass, but a very small
  amount of the volume that an atom occupies.
• 5. Lord Rutherford contributed to the
  understanding and concept of the nucleus as being
  a central core of the atom that was very dense
  and positively charged.

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Lord Rutherfords Experiment
Rutherford used a radioactive source that emitted
positively charged alpha particles and were
directed at the gold foil. The detector screen
was able to measure when and where the alpha
particles traveled.
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Lord Rutherfords Conclusion
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Description of Protons
1. Positively charged subatomic particle, having
an equal but opposite charge to an electron. 2.
Very (1840 times more) massive compared to an
electron but still very light in mass. Mass
1. 673 x 10-24 grams. 3. The number of protons
found in the nucleus, identifiesthe type of
atom. 4. Protons have a large, attractive force
for electrons, and weak attractive force for
neutrons. 5. Cannot be easily removed from the
nucleus, without the destruction of the atom.
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Description of Neutrons
1. Found in nucleus, and has a neutral charge (no
charge). 2. The number of neutrons can vary in
elements of the same type, causing differences in
masses.(ISOTOPES) 3. Very massive subatomic
particle compared to an electron, yet still very
light. Mass 1.675 x 10-24 grams 4.
Contributes with the proton to account for most
of the mass of an atom. 5. Cannot be easily
removed from the nucleus, without the destruction
of the atom.
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What are Isotopes?
1. ISOTOPES are atoms of the same element that
have different masses, due to the different
amounts of neutrons in the nucleus. 2. Some
isotopes are stable, while others are radioactive
and decay into other atoms. 3. Radioactive
isotopes sometimes are referred to as
nuclides. 4. Most isotopes are not given names,
we use the element's symbol and the mass number,
to identify it. 5. Hydrogens isotopes have been
named based on the number of nucleons in the
nucleus.
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The three isotopes of Hydrogen
Protium contains one proton, no neutrons and one
electron.
Deuterium contains one proton, one neutron, and
one electron.
Tritium contains one proton, two neutrons, and
one electron. Tritium is unstable and as a
result its nucleus breaks apart.
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Atomic Terms
Atomic number the number of an element that
designates the number of protons in the
nucleus. Atomic mass number the total number of
protons and neutrons in the nucleus of an atom.
Nuclide a general term to refer to any isotope
of any element. Z the letter or symbol used to
express the total number of protons or atomic
number of an element. A the letter or symbol
used to express the total number of protons and
neutrons or atomic mass number of an
element. Anode- electrode plate having a positive
charge. Cathode-electrode plate having a negative
charge.
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Using Atomic Number and Atomic Mass Number in
Chemistry
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Methods to express the Atomic Mass Unit

• Atomic mass number can be expressed as a relative
  mass unit, which is one that is compared to mass
  another element (carbon).
• Atomic mass number can be expressed as an average
  mass unit, which is one that is compared to mass
  another element (carbon).
• The periodic table expresses the atomic mass unit
  a weighted average mass of all of the naturally
  occurring isotopes.
• This is the reason that on the periodic table,
  the mass unit is rarely a whole number.

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Typical Ways of Writing Isotopes

• 238
• U
• 92

U-238
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RELATIVE ATOMIC MASS
1. One atom, CARBON was chosen as a standard for
comparing the masses of all other atoms. 2. Since
this type of scale of measurement is based on
comparison, this is called a relative unit. 3. By
definition One relative atomic mass unit (u) is
exactly 1/12 of the mass of a carbon-12 atom.
(amu or mu) 4. Today, we know the actual masses
of atoms, yet for most chemical calculations,
this is not needed.
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AVERAGE ATOMIC MASSES
1. Atomic masses found on the periodic table are
averages of naturally occurring mixtures of
isotopes. 2. The average mass is a weighted
average based on the frequency of occurrence in
nature of these isotopes. 3. This is not an
actual number that exists for an isotope it is a
calculated value, just like an average is a
calculated value.
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Example Average Mass Problem
(Mass number x decimal of occurrence) (mass
number x decimal of occurrence) the weighted
average atomic mass. O- 16 percentage of
occurrence ( 99.762 ) O-17 percentage of
occurrence (0.038) O- 18 percentage of
occurrence (0.200)
Solving this problem Oxygen 16
Oxygen 17 Oxygen 18 (16 x 0.99762) (17
x 0.00038) (18 x 0.0020) 15.962
0.00646 0.036
16.00446 amu
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Periodic Table Exercise
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Periodic Table Exercise
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19
21
19
K
Chromium
Cr
24
28
24
Uranium
U
92
92
92
Calcium
Ca
20
20
20
Silver
47
47
61
47
79
118
79
Au
Iron
26
26
56
26
Hg
200
80
80
Strontium
38
38
88
38
82
82
125
82
Pb
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Mole and Molar mass
TERMS Mole the amount of substance that
contains as many particles as there are atoms in
exactly 12 grams of carbon-12. Avogadros Number
represents 6.0221367 x 1023 particles in exactly
one mole of a pure substance. Molar mass the
mass of one mole of a pure substance. Practical
definitions One mole of an element is the number
of grams which is equal to the atomic mass of the
element found on the periodic table. Molar mass
is the gram amount of the substance that is equal
to one mole.
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Avogadros Number

• Avogadros Number represents 6.0221367 x 1023
  particles in exactly one mole of a pure
  substance.
• This is numerically
• 602 213 670 000 000 000 000 000 particles
• One mole of any pure substance always has 602
  213 670 000 000 000 000 000 particles.
• A mole is like the term one dozen.
• One dozen is equal to 12 of something.
• Do different substances have the same mass for
  one mole?
• NO each substance has its own unique atomic
  mass.

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Mole Problem
Formula Mole gram amount of a substance
atomic mass of the substance
If there are 240 grams of carbon, how many moles
does this represent? 1. Look up atomic mass of
carbon. (12.0 g/mol ) 2. Use formula Mole
gram amount of a substance atomic
mass of the substance Moles 240 g
12 g /mole Moles 20 moles of carbon
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Mass Problem
Formula Mass molar mass (atomic mass) x
moles
How many grams of aluminum, does 10 moles of
aluminum represent? 1. Look up atomic mass of
aluminum 2. Use formula Mass molar mass
(atomic mass) x moles Mass 27.0 g /mol
x 10 moles Mass 270 grams of aluminum