Unit 3:

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Unit 3 Atomic Structure
Chemistry Mr. Blake/Mr. Gower
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I. Atomic Structure
A. Section 4.1 Defining the Atom

• NOTE The Greek philosopher __________ (460 B.C.
  370 B.C.) was among the first to suggest the
  existence of atoms (from the Greek word atomos)
• He believed that atoms were ________ and
  _____________
• His ideas did agree with later scientific theory,
  but did not explain chemical behavior, and was
  ________________
• _______________ but just philosophy

Democritus
indivisible
indestructible
not based on the
scientific method
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1. Daltons Atomic Theory (experiment based!)
a. All elements are composed of tiny indivisible
particles called _____. b. Atoms of the same
element are _______. Atoms of any one element
are _______ from those of any other element.
atoms
identical
different
John Dalton (1766 1844)
c. Atoms of different elements combine in simple
whole-number ratios to form chemical _________ d.
In chemical reactions, atoms are combined,
separated, or rearranged but ____ changed into
atoms of another element.
compounds
never
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Problems with Daltons Atomic Theory?

• 1. matter is composed of indivisible particles
• Atoms Can Be Divided, but only in a nuclear
  reaction
• 2. all atoms of a particular element are
  identical
• Does Not Account for Isotopes (atoms of the same
  element but a different mass due to a different
  number of neutrons)! Different elements have
  different atoms. YES!
• 3. atoms combine in certain whole-number ratios
• YES! Called the Law of Definite Proportions
• 4. In a chemical reaction, atoms are merely
  rearranged to form new compounds they are not
  created, destroyed, or changed into atoms of any
  other elements.
• Yes, except for nuclear reactions that can
  change atoms of one element to a different
  element

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2. Sizing up the Atom
a. Elements are able to be subdivided into
smaller and smaller particles these are the
_____, and they still have __________ of that
element b. If you could line up 100,000,000
copper atoms in a single file, they would be
approximately ________ c. Despite their ________,
individual atoms ___ observable with instruments
such as scanning tunneling (electron) microscopes
atoms
properties
1 cm long
small size
are
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B. Section 4.2 Structure of the Nuclear Atom

• NOTE One change to Daltons atomic theory is
  that _______________ into subatomic particles
• NOTE ________________________ are examples of
  these fundamental particles
• NOTE There are many other types of particles,
  but we will study these three

atoms are divisible
Electrons, protons, and neutrons
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1. Discovery of the Electron
a. In 1897, J.J. Thomson used a _______ _______
to deduce the presence of a negatively charged
particle the _______
cathode
ray tube
electron
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Modern Cathode Ray Tubes
Television
Computer Monitor

• Cathode ray tubes pass electricity through a gas
  that is contained at a very low pressure.

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2. Mass of the Electron
Mass of the electron is 9.11 x 10-28 g
The oil drop apparatus
Robert Millikan
a. 1916 _____________ determines the _____ of
the electron 1/1840 the mass of a hydrogen atom
has one unit of negative charge
mass
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3. Conclusions from the Study of the Electron
a. Cathode rays have identical properties
regardless of the element used to produce them.
All elements must contain identically charged
electrons. b. Atoms are neutral, so there must be
______________ in the atom to balance the
negative charge of the electrons c.
_______________________ that atoms must contain
other particles that account for most of the mass
positive particles
Electron have so little mass
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Eugen Goldstein
d. ____________ in 1886 observed what is now
called the ______ - particles with a positive
charge, and a relative mass of 1 (or 1840 times
that of an electron) e. 1932 _____________
confirmed the existence of the ______ a
particle with ________, but a mass nearly ____ to
a proton
proton
James Chadwick
neutron
no charge
equal
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4. Subatomic Particles
Particle Charge Mass (g) Location
Electron (e-) -1 9.11 x 10-28 Electron cloud
Proton (p) 1 1.67 x 10-24 Nucleus
Neutron (no) 0 1.67 x 10-24 Nucleus
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5. Thomsons Atomic Model
J. J. Thomson
a. Thomson believed that the ________ were like
plums embedded in a positively charged pudding,
thus it was called the __________ model.
electrons
plum pudding
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6. Ernest RutherfordsGold Foil Experiment - 1911
a. Alpha particles are helium nuclei - The alpha
particles were fired at a thin sheet of gold
foil b. Particles that hit on the detecting
screen (film) are recorded
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7. Rutherfords problem
a. In the following pictures, there is a target
hidden by a cloud. To figure out the shape of the
target, we shot some beams into the cloud and
recorded where the beams came out. Can you figure
out the shape of the target?
Target 2
Target 1
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b. The Answers
Target 2
Target 1
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8. Rutherfords Findings

• Most of the particles passed right through
• A few particles were deflected
• c. VERY FEW were greatly deflected

Like howitzer shells bouncing off of tissue
paper!
d. Conclusions
small
1) The nucleus is _____ 2) The nucleus is
_____ 3) The nucleus is _______ charged
dense
positively
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9. The Rutherford Atomic Model

• a. Based on his experimental evidence
• 1) The atom is mostly empty space
• 2) All the positive charge, and almost all the
  mass is concentrated in a small area in the
  center. He called this a ______
• 3) The nucleus is composed of ______ and ________
  (they make the nucleus!)
• 4) The electrons distributed around the nucleus,
  and occupy most of the ______
• 5) His model was called a ___________

nucleus
protons
neutrons
volume
nuclear model
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1. Atomic Number
C. Section 4.3 Distinguishing Among Atoms
identical

• a. Atoms are composed of _______ protons,
  neutrons, and electrons
• b. How then are atoms of one element different
  from another element?
• c. Elements are different because they contain
  different numbers of ________
• d. The ____________ of an element is the
  _______________ in the nucleus
• e. ___________________________

PROTONS
atomic number
number of protons
protons in an atom electrons
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2. Definition Atomic number (Z) of an element is
the number of protons in the nucleus of each atom
of that element.
Element of protons Atomic (Z)
Carbon
Phosphorus
Gold
6
6
15
15
79
79
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3. Mass Number
Definition Mass number is the number of protons
and neutrons in the nucleus of an isotope
Mass p n0
Nuclide p n0 e- Mass
Oxygen - 10
- 33 42
- 31 15
18
8
18
8
75
75
Arsenic
33
31
16
Phosphorus
15
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4. Nuclear/Complete Symbols

• a. Contain the symbol of the element, the mass
  number and the atomic number.

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5. Symbols
a. Find each of these 1) number of protons 2)
number of neutrons 3) number of electrons 4)
Atomic number 5) Mass Number
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45
35
35
80
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• b. If an element has an atomic number of 34 and
  a mass number of 78, what is the
• number of protons
• number of neutrons
• number of electrons
• complete symbol

34
44
34
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• d. If an element has 78 electrons and 117
  neutrons what is the
• Atomic number
• Mass number
• 3) number of protons
• 4) complete symbol

78
195
78
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6. Isotopes

• a. Dalton was wrong about all elements of the
  same type being identical
• b. Atoms of the same element can have different
  numbers of _______.
• c. Thus, different mass numbers.
• d. These are called _______.

neutrons
isotopes
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Isotopes
Frederick Soddy

• e. _____________(1877-1956) proposed the idea of
  isotopes in 1912
• f. _______ are atoms of the ____ ______ having
  different masses, due to varying numbers of
  neutrons.
• g. Soddy won the Nobel Prize in Chemistry in 1921
  for his work with isotopes and radioactive
  materials.

same
Isotopes
element
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7. Naming Isotopes

• a. We can also put the mass number after the name
  of the element
• b. Examples
• carbon-12
• carbon-14
• uranium-235

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same element
Isotopes
c. _______ are atoms of the ___________ having
________ masses, due to varying numbers of
neutrons.
different
Isotope Protons Electrons Neutrons Nucleus
Hydrogen1 (protium) 1 1 0
Hydrogen-2 (deuterium) 1 1 1
Hydrogen-3 (tritium) 1 1 2
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8. Isotopes
a. Elements occur in nature as _______ of _______.
mixtures
isotopes
b. Isotopes are atoms of the same element that
differ in the _______ _______.
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9. IONS

• a. ____ are atoms or groups of atoms with a
  positive or negative charge.
• b. _________ an electron from an atom gives a
  _____ with a ____________
• c. ______ an electron to an atom gives an _____
  with a ____________.
• d. To tell the difference between an atom and an
  ion, look to see if there is a charge in the
  _________! Examples Na Ca2 I- O-2
• Na Ca I O

Ions
Taking away
cation
positive charge
Adding
anion
negative charge
superscript
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f. An anion forms when an atom gains one or more
electrons
e. A cation forms when an atom loses one or more
electrons.
F e- --gt F-
Mg --gt Mg2 2 e-
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NOTE In General

• metals (Mg) lose electrons ---gt cations
• nonmetals (F) gain electrons ---gt anions

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Learning Check Counting

• State the number of protons, neutrons, and
  electrons in each of these ions.
• 39 K 16O -2 41Ca 2
• 19 8 20
• p ______ ______ _______
• no ______ ______ _______
• e- ______ ______ _______

20
19
8
8
21
20
18
10
18
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One Last Learning Check

• Write the nuclear symbol form for the following
  atoms or ions
• A. 8 p, 8 n, 8 e- ___________
• B. 17p, 20n, 17e- ___________
• C. 47p, 60 n, 46 e- ___________

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Charges on Common Ions
By losing or gaining e-, atom has same number of
e-s as nearest Group 8A atom.
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Example A student has a test percentage of
78 a lab percentage of 92 and has completed
homework at 100. Her weighted average grade is
computed as (78 X 0.6) (92 X 0.20) (100
X 0.20) 84.5
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10. Atomic Mass

• a. How heavy is an atom of oxygen?
• It depends, because there are different _____ of
  oxygen atoms.
• b. We are more concerned with the
• _________________.
• c. This is based on the abundance (percentage) of
  each variety of that element in nature.
• d. We dont use grams for this mass because the
  numbers would be too small.

kinds
average atomic mass
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11. Measuring Atomic Mass

• a. Instead of grams, the unit we use is the
  ______________ (amu)
• b. It is defined as one-twelfth the mass of a
  carbon-12 atom.
• c. Carbon-12 chosen because of its _____
• ______.
• d. Each isotope has its own atomic mass, thus we
  determine the average from percent abundance.

Atomic Mass Unit
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12. To calculate the average

• a. Multiply the atomic mass of each isotope by
  its abundance (expressed as a decimal), then add
  the results.
• b. If not told otherwise, the mass of the isotope
  is expressed in _____________ (amu)

atomic mass units
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13. Atomic Masses
Atomic mass is the average of all the naturally
occurring isotopes of that element.
Isotope Symbol Composition of the nucleus in nature
Carbon-12 12C 6 protons 6 neutrons 98.89
Carbon-13 13C 6 protons 7 neutrons 1.11
Carbon-14 14C 6 protons 8 neutrons lt0.01
Carbon 12.01
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- Page 117
Question
Knowns and Unknown
Solution
Answer
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• 14. Calculate the atomic mass of carbon.
• a. Isotopes abundance Atomic mass
• Carbon-12 98.89 12.000 amu
• Carbon-13 1.11 13.003 amu
• .
• b. Lithium has two isotopes. If lithium-6 has a
  mass of 6.015 and 7.42 occurrence, what is the
  abundance and mass of lithium -7?

Atomic mass
() (mass)
() (mass)
..
Atomic mass
(0.9889) (12.000)
(0.0111) (13.003)
11.87 .144 12.01 amu
6.941
(0.0742)(6.015)
(0.9258)(x)
7.015 amu
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15. The Periodic Table A Preview
a. A periodic table is an arrangement of
elements in which the elements are separated into
groups based on a set of repeating properties 1)
The periodic table allows you to easily compare
the properties of one element to another
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b. Each horizontal row (there are 7 of them) is
called a _____ c. Each vertical column is called
a ____________ 1) Elements in a _____ have
similar chemical and physical properties 2)
Identified with a number and either an A or B
period
group or family
group
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II. The Periodic Table
A. Section 6.1 Organizing the Elements
NOTES

• A few elements, such as gold and copper, have
  been known for thousands of years - since ancient
  times
• Yet, only about __ had been identified by the
  year 1700.
• As more were discovered, chemists realized they
  needed a way to ________ the elements.

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organize
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properties

• Chemists used the _________ of elements to sort
  them into groups.
• In 1829 J. W. Dobereiner arranged elements into
  _____ groups of three elements with similar
  properties
• One element in each triad had properties
  intermediate of the other two elements

triads
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1. Mendeleevs Periodic Table

• a. By the mid-1800s, about 70 elements were known
  to exist
• b. Dmitri _________ a Russian chemist and
  teacher
• c. Arranged elements in order of
• _________________
• d. Thus, the first Periodic Table

Mendeleev
increasing atomic mass
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2. Mendeleev
He left blanks

• a. ___________ for yet undiscovered elements
• b. When they were discovered, he had made good
  predictions
• c. But, there were problems
• Such as Co and Ni Ar and K Te and I

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3. A Better Arrangement
Moseley

• a. In 1913, Henry ______ British physicist,
  arranged elements according to increasing
• ____________
• b. The arrangement used today
• c. The symbol, atomic number mass are basic
  items included-textbook page 162 and 163

atomic number
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4. The Periodic Law

• a. When elements are arranged in order of
  increasing atomic number, there is a periodic
  repetition of their physical and chemical
  properties.
• b. Horizontal rows ______
• 1) There are __ periods
• c. Vertical column _____ (or family)
• 1) Similar physical chemical prop.
• 2) Identified by number letter (IA, IIA)

periods
7
group
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5. Areas of the Periodic Table

• Three classes of elements are 1) _____,
  2) ________, and 3) _________
• Metals _______ conductors, have luster, ductile,
  malleable
• Nonmetals generally brittle and non-lustrous,
  poor conductors of ____ and electricity

metals
nonmetals
metalloids
electrical
heat
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Areas of the Periodic Table
gases

• Some nonmetals are _____ (O, N, Cl) some are
  brittle solids (S) one is a fuming dark red
  liquid (Br)
• Notice the heavy, stair-step line?
• _________ border the line-2 sides
• Properties are __________ between metals and
  nonmetals

Metalloids
intermediate
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Squares in the Periodic Table
symbols

• The periodic table displays the ______ and _____
  of the elements, along with information about the
  structure of their atoms
• Atomic ______ and atomic _____
• Black symbol solid red gas ____ _____
  (from the Periodic Table on our classroom wall)

names
mass
number
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Groups of Elements - Family Names
alkali metals

• Group IA (1) __________
• Forms a base (or alkali) when _______ with
  water (not just dissolved!)
• Group 2A (2) ________________
• Also form bases with water do not dissolve well,
  hence earth metals
• Group 7A (17) _______
• Means salt-forming
• Group 8A (18) _________
• Nonreactive because of their electron
  configuration

reacting
alkaline earth metals
halogens
noble gases
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ELEMENTS THAT EXIST AS DIATOMIC MOLECULES
Remember HOFBrINCl These elements only exist as
PAIRS. Note that when they combine to make
compounds, they are no longer elements so they
are no longer in pairs!
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Nuclear Chemistry
Chemistry Unit 4 Chapter 25
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Mass Defect

• Difference between the mass of an atom and the
  mass of its individual particles.

4.00260 amu
4.03298 amu
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Nuclear Binding Energy

• Energy released when a nucleus is formed from
  nucleons.
• High binding energy stable nucleus.

E mc2
E energy (J) m mass defect (kg) c speed of
light (3.00108 m/s)
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Nuclear Binding Energy
Unstable nuclides are radioactive and undergo
radioactive decay.
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Types of Radiation

• Alpha particle (?)
• helium nucleus

paper
2

• Beta particle (?-)
• electron

1-
lead

• Positron (?)
• positron

1
concrete

• Gamma (?)
• high-energy photon

0
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Nuclear Decay

• Alpha Emission

Numbers must balance!!
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Nuclear Decay

• Beta Emission

• Positron Emission

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Nuclear Decay

• Electron Capture

• Gamma Emission
• Usually follows other types of decay.
• Transmutation
• One element becomes another.

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IQ 1

• Balance the following equations

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Nuclear Decay

• Why nuclides decay
• need stable ratio of neutrons to protons

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Belt of Stability and Radioactive Decay
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Half-life

• Half-life (t½)
• Time required for half the atoms of a radioactive
  nuclide to decay.
• Shorter half-life less stable.

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Half-life
mf final mass mi initial mass n of half-lives
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Half-life

• Fluorine-21 has a half-life of 5.0 seconds. If
  you start with 25 g of fluorine-21, how many
  grams would remain after 60.0 s?

WORK mf mi (½)n mf (25 g)(0.5)12 mf 0.0061
g
GIVEN t½ 5.0 s mi 25 g mf ? total time
60.0 s n 60.0s 5.0s 12
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• Example How much of a 500. g sample of
  Uranium-235 would be left after five half-lives?

Mi 500 g
n 5
(n of half-lives)
Mf ?
mf mi (½)n mf (500 g)(0.5)5 mf 15.6 g
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• Example A 16.00 mg sample of Radon-222 decays
  to 0.250 mg after 24 hours. Determine the
  half-life.

16? 8 ? 4 ? 2 ? 1 ? 0.5 ? 0.250 6 half lives
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• Example The half-life of molybdenum-99 is 67
  hours. How much of a 1.000 mg sample is left
  after 335 hours?

Mi 1.000 mg
mf mi (½)n mf (1.000 mg)(0.5)5 mf 0.03125 mg
Half-life 67 h
Rxn time 335 h
Mf ?
n 335 / 67 5
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Learning Check!

• The half life of I-123 is 13 hr. How much of a
  64 mg sample of I-123 is left after 39 hours?

mf mi (½)n mf (64 mg)(0.5)3 mf 8.0 mg
Mi 64 mg
n 3
Mf ?
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• Half-life Lab
• Procedure
• 1. Each lab group will acquire a sample of 50
  pennies in a cup.
• 2. Count pennies to make sure you have 50
  pennies.
• Enter 50 in Shake 0 row for Trial 1, 2, 3
  and 150 for ? (Sum of) of trials.
• 4. Shake the cup of pennies. Pour the pennies
  on to the lab bench.
• 5. Remove all pennies that land on heads.
  They have decayed.
• 6. Count only the remaining pennies (the pennies
  that landed on tails). Record data.
• 7. Place only the remaining pennies (tails)
  into the cup and shake again. Repeat steps 4-7
  until all pennies have decayed.
• 8. Repeat the process two more times and record
  data under Trial 2 3.

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• Data Collect data for three trials in the table.
• Data Analysis Prepare a graph to represent the
  decay of your sample (? of trials (y-axis) vs.
  Shake (x-axis))
• Prepare a graph in your lab book
• Graph the of undecayed atoms (? of trials)
  (y-axis) versus the Shake (x-axis). Label the
  x and y axes, including units (if applicable).
  Make graph large (at least 2/3 pg.). Draw a best
  fit curve that represents your data. Use a
  Ruler!
• Plot the Shake for the ? of trials using the
  best fit curve.
• Determine the half life of your sample in terms
  of of shakes using your graph.

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Graphing the Results

• Important !!
• Title every graph and label each axis (include
  units)
• Graphs should be at least 2/3 page
• Use a ruler
• Circle all data points
• Use a best-fit line (no connect the dots!)
• Find the average half-life (in of trials) of
  your sample by interpolating your curve at
  exactly 75, 37.5, and 18.75 pennies undecayed)

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Half-Life Lab

Title
At least 2/3 of pg
Use Ruler for axis
Label Axis
Circle Data Points
Best Fit Curve
S of Trials
Convenient s

1
1
1.2
3.2
1.1 shake

3

Half-life 1.1 shake
0
Shake
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F ission

• splitting a nucleus into two or more smaller
  nuclei
• 1 g of 235U 3 tons of coal

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F ission

• chain reaction - self-propagating reaction
• critical mass - mass required to sustain a
  chain reaction

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Fusion

• combining of two nuclei to form one nucleus of
  larger mass
• thermonuclear reaction requires temp of
  40,000,000 K to sustain
• 1 g of fusion fuel 20 tons of coal
• occurs naturally in stars

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Fission vs. Fusion
FISSION
FUSION

• 235U is limited
• danger of meltdown
• toxic waste
• thermal pollution

• fuel is abundant
• no danger of meltdown
• no toxic waste
• not yet sustainable

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Nuclear Power
Cooling Tower

• Fission Reactors

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Nuclear Power

• Fission Reactors

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Nuclear Power

• Fusion Reactors (not yet sustainable)

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Nuclear Power

• Fusion Reactors (not yet sustainable)

National Spherical Torus Experiment
Tokamak Fusion Test Reactor Princeton University
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Synthetic Elements

• Transuranium Elements
• elements with atomic s above 92
• synthetically produced in nuclear reactors and
  accelerators
• most decay very rapidly

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Radioactive Dating

• half-life measurements of radioactive elements
  are used to determine the age of an object
• decay rate indicates amount of radioactive
  material
• EX 14C - up to 40,000 years 238U and 40K - over
  300,000 years

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Nuclear Medicine

• Radioisotope Tracers
• absorbed by specific organs and used to diagnose
  diseases
• Radiation Treatment
• larger doses are used to kill cancerous cells
  in targeted organs
• internal or external radiation source

Radiation treatment using ?-rays from cobalt-60.
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Nuclear Weapons

• Atomic Bomb
• chemical explosion is used to form a critical
  mass of 235U or 239Pu
• fission develops into an uncontrolled chain
  reaction
• Hydrogen Bomb
• chemical explosion ? fission ? fusion
• fusion increases the fission rate
• more powerful than the atomic bomb

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Others

• Food Irradiation
• ? radiation is used to kill bacteria
• Radioactive Tracers
• explore chemical pathways
• trace water flow
• study plant growth, photosynthesis
• Consumer Products
• ionizing smoke detectors - 241Am