1. Nuclear Chemistry Brief history of nuclear related discoveries
Electron, proton, neutron
Nuclear transformations
Natural radioactivity
Half Life, carbon dating
Nuclear chemistry equations
Chain reaction, atom bomb
Applications
Nuclear reactors
Radioisotopes
Personl Exposure
Radon, other natural sources
3-mile Island (USA), Chernobyl (USSR), Japan
Potassium Iodide protection

2. Atomic versus Nuclear What’s the Difference?
Atomic properties are those of an atom
Chemical reactions (gain/loss of electrons)
Emission of light (electron orbit jumps)
Bonds between elements (covalent, ionic)
Nuclear properties are within the atom
Construction of nucleus (electrons, protons)
Radioactive disintegration
Fission/fusion of nuclei, element conversion

3. Atomic number Atomic Number (and element number) Atomic Mass Number
Number protons = number electrons = “Z”
Atomic Mass Number
Total Number of Protons defining the element = Z
Total number of Neutrons in element nucleus = N
Total mass of nucleus = A = Z + N
Electron mass is ignored 1/1836 =0.054% (negligible)
Isotopes
Same atomic number, different number of neutrons
Large variations of isotopes between elements
Isotope significance (e.g. U-235 vs U-238, C-14 vs C-12)
Atomic Weight (most often used)
Weighted average of isotope masses
What’s on the periodic chart

4. Weighted Atomic Mass (natural values)
Lots of isotopes of Uranium exist, all different masses
Most prevalent isotope is U-238 at %
Most useful isotope is U-235 at 0.711%
Other isotopes rather rare, relatively insignificant
“Weighted Average” recognizes quantity
Sum of mass contributions = the weighted average value
This is what we see on the periodic charts.

5. Nuclear Changes
Nuclear reaction: A reaction that changes an atomic nucleus, usually causing the change of one element into another. A chemical reaction never changes the nucleus.
Different isotopes of an element have essentially the same behavior in chemical reactions but often have completely different behavior in nuclear reactions.
The rate of a nuclear reaction is unaffected by a change in temperature or pressure (within the range found on earth) or by the addition of a catalyst.
The nuclear reaction of an atom is essentially the same whether it is in a chemical compound or in an uncombined, elemental form.
The energy change accompanying a nuclear reaction can be up to several million times greater than that accompanying a chemical reaction.

6. Nuclear Nomenclature
The atomic number, written below and to the left of the element symbol, gives the number of protons in the nucleus and identifies the element.
The mass number, written above and to the left of the element symbol, gives the total number of nucleons, a general term for both protons (p) and neutrons (n).
The most common isotope of carbon, for example, has 12 nucleons: 6 protons and 6 neutrons:

7. Writing Nuclear Reactions
“stacked” numbers difficult to write
Possible but difficult using a word processor
Alternative is “front and back” values
Carbon 14 = 6C14
Front value is atomic number Z
Back value is atomic mass A

8. 60 years of discoveries
Chemistry theories, 56 theses, Arrhenius
Proton Discovery, Goldstein
1895 – X-Ray discovery, Roentgen
Radioactivity Discovery, Baquerel
Electron Discovery, J.J. Thompson
Radioactive Element separation, Curie
Equivalence of mass & energy E=mC2, relativity, photoelectric effect, A. Einstein
general relativity, proven in 1919, A. Einstein
Neutron Discovery, Chadwick
Nuclear chain reaction proposed, Szilard
Nuclear fission discovered
First operational nuclear reactor, Fermi
First warfare use of nuclear energy

9. Cathode Rays (electrons) first demonstrated by Crookes in1895 Early investigator of radiation inside electrical discharge tubes, eventually leading to CRT (Television) tubes. He was one of the first to experiment with radioactivity and its ability to make certain minerals glow. He also invented the ”radiometer” still in use as an educational toy.

10. ELECTRON - J.J. Thompson 1897 found a new particle “boiling off” a heated filament which had <1/1000 mass of hydrogen. It had a negative charge by its magnetic and/or electrostatic deflection. Using similar apparatus he discovered isotopes of the same element with different mass, which led to science of mass spectrometry

11. PROTON, Goldstein in1888 Used high voltage to ionize gases, accelerating particles through holes in cathode, causing “canal rays” (trails looked like canals). Particles were positive. Hydrogen particles later identified as protons by Rutherford in 1919

12. NEUTRON, Chadwick 1932 observed a new form of penetrating radiation, which had no charge (not protons or electrons)

13. NEUTRINO Predicted by Wolfgang Pauli in 1930, based on conservation discrepancies. the “little neutron” Indirectly observed in 1942 and1946 via interactions with other particles, directly observed in 1972 “bubble chamber”

14. Discovery and Nature of Radioactivity
In 1896, the French physicist Henri Becquerel noticed a uranium-containing mineral exposed a photographic plate that had been wrapped in paper.
Marie and Pierre Curie investigated this new phenomenon, which they termed radioactivity: The spontaneous emission of radiation from a nucleus.
Ernest Rutherford established that there were at least two types of radiation, which he named alpha and beta. Shortly thereafter, a third type of radiation was found and named for the third Greek letter, gamma.

15. Baquerel – observed radioactivity 1896 Photographic plate accidentally exposed by Uranium Baquerel is SI unit of radiation, Bq = disintegrations/sec 1 Curie = radiation from 1 gram of Radium = 3.7*10^10 Bq

16. Radioactivity Emissions from Pitchblende (uranium ore)
Found to expose photographic film
We will measure pitchblende today
3 common types of nuclear radiation
Alpha (α), helium nuclei particle, 2He4 = 2p+2n
Very strong but not very penetrating
Beta (β), an electron particle, -1e0
Mildly penetrating, stopped by thick paper
Gamma (γ), radiation similar to X-Ray
Very penetrating, used for imaging 16

17. Stable and Unstable Isotopes
Every element in the periodic table has at least one radioactive isotope, or radioisotope, and more than 3300 radioactive isotopes are known.
Their radioactivity is the result of having unstable nuclei. Radiation is emitted when an unstable radioactive nucleus, or radionuclide, spontaneously changes into a more stable one.
There are only 264 stable isotopes among all the elements.
All isotopes of elements with atomic numbers higher than that of bismuth (83) are radioactive.

18. Radioactivity Results from unstable elements
Heavy elements formed inside stars
Formed and stable at extreme temperatures
Unstable and disintegrate at earth temperatures
Half-Life
Time it takes for ½ of material to disintegrate
Uranium 238 is 4.5 billion years, same as earth’s age
There was twice as much uranium when earth formed
A non-linear scale (e.g ½ of ½, etc)
A natural nuclear reactor happened in Okla, Africa

19. Marie & Pierre Curie, 1905 Separated tons of mineral pitchblende to discover and isolate Radium, and polonium (named for Poland). The standard measure for radioactivity is the Curie = Ci 19

20. For elements in the first few rows of the periodic table, stability is associated with a roughly equal number of neutrons and protons.
As elements get heavier, the number of neutrons relative to protons in stable nuclei increases.
Lead-208, for example, the most abundant stable isotope of lead, has 126 neutrons and 82 protons in its nuclei.

21. Alpha rays move about ~0.1c and can be stopped by a few sheets of paper or by the top layer of skin.
Beta rays move at up to 0.9c and have about 100 times the penetrating power of a particles. A block of wood or heavy clothing is necessary to stop b rays.
Gamma rays move at c and have about 1000 times the penetrating power of a rays. A lead block several inches thick is needed to stop g rays.

22. When passed between two charged plates:
Alpha rays, helium nuclei (He+2 ), bend toward the negative plate because they have a positive charge.
Beta rays, electrons (e- ), bend toward the positive plate because they have a negative charge.
Gamma rays, photons (g), do not bend toward either plate because they have no charge.

23. Nuclear Decay
Nuclear decay: The spontaneous emission of a particle from an unstable nucleus.
Transmutation: The change of one element into another.
The equation for a nuclear reaction is not balanced in the usual chemical sense because the kinds of atoms are not the same on both sides of the arrow. A nuclear equation is balanced when the number of nucleons and the sums of the charges are the same on both sides.

24. Nuclear Decay Decrease in atomic number Increase in atomic number
Loss of alpha particle with positive charge
Mass of 2 protons and 2 neutrons
Loss of 2 protons reduces element # by 2
Uranium 92U238 into Thorium 90Th He4
Increase in atomic number
Loss of electron with negative charge
Charge of nucleus increases by 1
Next higher element formed
Thorium 90Th234  Proactinium 91Pa e0
Proactinium 91Pa234  Uranium 92U e0

25. During alpha emission, the nucleus loses two protons and two neutrons.
Emission of an a particle from an atom of uranium-238 produces an atom of thorium-234.

26. Beta emission involves the decomposition of a neutron to yield an electron and a proton.
Iodine-131, a radioisotope used in detecting thyroid problems, undergoes nuclear decay by b emission to yield xenon-131.

27. Anti-Particles
In 1928 Paul Dirac predicted that all particles should have opposites called anti-particles.
The first of these was discovered in 1932 by Carl Anderson. This was an electron with a positive electric charge (+1). This particle is the anti-electron (also called a positron). It is identical in every respect to the electron apart from its electric charge.
When an electron and positron come into contact, they mutually annihilate each other producing a flood of energy in accordance with Einstein's famous equation, E=mC2

28. Anti-matter
Anti-particles making up atoms are same as conventional particles except for charge
Electron (e-)  Positron (e+), LNL
Proton (p+)  Anti-Proton (p-), UCB
Hydrogen  Anti-Hydrogen; CERN
Matter+Antimatter  pure energy
Antimatter engines in Star Trek !
Parts of universe could be “anti-matter”
Same properties, light emission, etc.
Matter + Antimatter  pure energy
Something to avoid ….

29. Positron emission involves the conversion of a proton in the nucleus into a neutron plus an ejected positron.
A positron has the same mass as an electron but a positive charge.
Potassium-40 undergoes positron emission to yield argon-40, this hppens in your body which contains tiny amounts of K-40.

30. Radioactive Half-Life
Rates of nuclear decay are measured in units of half-life , defined as the amount of time required for one-half of the radioactive sample to decay.
Each passage of a half-life causes the decay of one half of whatever sample remains. The half-life is the same no matter what the size of the sample, the temperature, or any other external conditions.

31. All nuclear decays follow the same curve, 50% of the sample remains after one half-life, 25% after two half-lives, 12.5% after three half-lives, and so on.

32. Carbon-14 Carbon14 discovered 1940 at UC Berkeley
Formed in upper atmosphere from nitrogen
Solar neutrons convert nitrogen to carbon-14
Carbon-14 becomes carbon dioxide
Carbon-14 dioxide absorbed by plant life
Animal life eats the plant life, absorbs C-14
Steady state evolves … until object dies
After death, the decay rate reduces C-14

34. Decay of Carbon-14 Neutron turns into proton + electron (beta particle) Mass remains at 14, but carbon becomes nitrogen intensity of electron emission indicates object’s age

35. 14C is continually produced in the atmosphere and incorporated into life cycles, so 14C amount in living things is constant. Upon death, no more 14C is absorbed, so concentration decreases. Measuring the remaining radioactivity provides an age estimate. The half life of 14C is 5730 years. The method is good for estimating age of objects 500 to 50,000 years old. 35

36. Carbon dating methodology counts/min versus age of sample, assumes C-14 formation at a constant rate

37. Half Life calculations
Reduction in material is non-linear
original½ at half-life¼ at 2 half-lives
General equation:
n1/n0 = (0.5)^(#of half lives)
n1/n0 = (0.5)^1  50% remaining
n1/n0 = (0.5)^2  25% remaining
n1/n0 = (0.5)^3  12.5% remaining

38. Half Life calculations
Carbon dating example: (reverse example)
old object has 12.79% normal C14
Half life of C14 is 5730 years
n1/n0 = (0.5)^(number of half lives)
Take natural logs to eliminate exponents:
Ln(n1/n0)= Ln(0.5)*(t / 5730)
Ln(0.1279/1.00)= Ln(0.5)*(t / 5730)
Ln(0.1279)= (t / 5730)
= * (t / 5730)
t = ( )*(5730 years)/ = 17,002 years

39. K-40 is part of your body K-40 is the largest radiation source in you
Human body has ≈160 grams of Potassium
0.0117% K-40*160 grams  g K-40
This produces 4,400 disintegrations/sec
Emission mostly β (beta ray or electrons)
We will look at a potassium sample today

40. Radioisotopes used internally for medical applications will have short half-lives so that they decay rapidly and do not remain in the body for prolonged periods (Tc-99, ½=6hrs)

42. Ionizing Radiation
A large dose of ionizing radiation can destroy living cells, causing death.
A small dose of ionizing radiation may not cause visible symptoms but might lead to a genetic mutation or cancer.

43. What’s an “MeV”? Alternative energy unit used in physics
Equals 1 electron passing through 1 volt field
Same as 1.6*10-19 Joule
Convenience is simple numbers, fewer zeros
Energy of visible light in eV in diagram below
Nuclear particles & rays >106 times light energy
1 MeV = 10^6 eV

44. Light energy in electron volts

45. Radiation Intensity
Health professionals who work with X rays or other kinds of ionizing radiation protect themselves by surrounding the source with a thick layer of lead or other dense material.
Protection is also afforded by controlling the distance between the worker and the radiation source because radiation intensity (I) decreases with the square of the distance from the source. We will demonstrate inverse square law this week in lab.
The intensities (I) of radiation at two different distances (d) are given by the equation:
I1d12 = I2d22

46. Inverse Square Law

47. Inverse square law Intensity of radiation falls with distance
A non-linear relationship
Twice as far  ¼ the intensity
Works with all forms of point source radiation
Why planet Mercury is so hot, Venus so cold

48. Inverse Square Example
Radiation intensity is 2250 at 2 meters, what is intensity at 3 meters?
I1 * d12 = I2 * d22
2250 * 22 = I2 * 32
I2 = 2250*4/9 = 1000
Moving away by 1 meter cuts radiation by more than half

49. Mass into Energy Enormous ratio between mass & Energy
c = 3*10^8 meters/sec
c2 = 9*10^16 meters2/sec2
How much energy is that ?
1 gram U235 converts to 3.4*10^8 kcal
Hiroshima bomb converted only a few grams 49

50. Equivalence of Mass & Energy Albert Einstein’s famous equation E=mc250

51. E=mc2, so light has mass E=mc2, E= hc/λ, so m=h/(c. λ)= 3
E=mc2, so light has mass E=mc2, E= hc/λ, so m=h/(c*λ)= 3.4E-36 kg/photon This is about 1 millionth the mass of an electron (Although very tiny, mass of light has important consequences 51

52. Space distortion due to mass General Relativity predicted light deflection by mass in 1916, proven by experiment in 1919 52

53. Gravity Well or “Black Hole” space is so distorted that light cannot pass nearby without falling in … a consequence of general relativity 53

54. Chain Reaction Postulated in 1930’s
Fission induces more fission, an avalanche
“critical mass” required to sustain reaction
Only certain isotopes sustain the reaction
Uranium 235 and Plutonium 239 most common
Show You-Tube simulation
First atomic reactor University of Chicago
Concentrated U235 required
The critical mass for lower-grade uranium depends strongly on the grade:
20% U235 requires over 400 kg
15% U235 requires over 600 kg.

55. Critical Mass
A chain reaction requires a minimum amount of material for to be self sustaining
Plutonium-239 the size of a softball is enough to make a nuclear weapon.

56. U235 %
Natural 0.711% U235 cannot sustain a chain reaction. Too much distance & shielding between fissionable atoms.
Minimal enrichment OK for reactors
High enrichment for weapons, reduces minimum mass and increases reaction rate.

57. First Nuclear Fission Bomb

58. Nuclear Fusion Nuclear fusion: The joining together of light nuclei.
Light nuclei such as the isotopes of hydrogen release enormous amounts of energy when they undergo fusion. It is fusion reactions of hydrogen nuclei to produce helium that powers our sun and other stars.
The necessary conditions for nuclear fusion are not easily created on earth.
In stars the temperature is on the order of 2 x 107 K and pressures approach 1 x 105 atmospheres. At these extremes, nuclei are stripped of all their electrons and have enough kinetic energy that nuclear fusion readily occurs.

59. Fusion is primary energy source
Universe is mostly Hydrogen
75% by mass, 90% of all atoms
Hydrogen fuses at high temperature to Helium
Helium mass less that Hydrogen starting material
Balance of mass converted to energy
Fusion reaction powers the sun
Sun is a continuously operating Hydrogen bomb
Sun is also our primary light source
Fundamental basis for life on earth
Numerous energy conversion processes

60. Nuclear Fusion in the Sun

61. Nuclear Fusion Principle source of energy in stars from fusing hydrogen atom of helium weighs slightly less than 4 hydrogen, balance is energy

62. Heavier elements formed in stars formation due extreme temperatures & pressures inside star

63. Nuclear Fusion (Hydrogen) Bomb

64. Medical Issues Radiation poisoning Radiation protection
Bad if uncontrolled fallout (Chernobyl)
Good if controlled attack on tumors
Same idea in chemotherapy, selective poison
Radiation protection
Radioactive iodine from nuclear accidents
Thyroid collects & concentrates iodine
Use of potassium iodide to saturate thyroid prevents uptake of radioactive fallout iodine

65. Thyroid Cancer

66. Polonium Poisonings Alexander Litvenenko ex-KGB agent and critic of Russian politics. At lunch in London with ”friends” he ingested 10 micrograms Po-210, 200x a lethal dose of extremely radioactive alpha-emitter, and died 3 weeks later. There is some suspicion Yasser Arafat may have been similarly poisoned

67. Polonium-210 poisoning

68. Chernobyl reactor meltdown 2.3% radiation increase world-wide

69. Chernoble fallout victims

70. Detecting Radiation
We cannot see, hear, smell, touch, or taste radiation, no matter how high the dose. We can, however, detect radiation by taking advantage of its ionizing properties.
The simplest device for detecting exposure to radiation is the photographic film badge.
The film is protected from exposure to light, but any other radiation striking the badge causes the film to fog. 70

71. The Geiger counter is an argon-filled tube
The Geiger counter is an argon-filled tube. The inner walls are given a negative charge, and a wire in the center is given a positive charge.
Radiation ionizes the argon atoms, which briefly conduct a current between the walls and the center electrode.
The current is detected, amplified, and used to produce a clicking sound or to register on a meter. 71

72. When Ionizing radiation strikes certain crystals or plastics, a small flash of light results, which can be measured

73. Scintillation, emission of light Observed by Crooks in 1895, using luminescence of ZnS, which became basis of “picture tubes” in television. Radiation generating a single photon of light can be amplified108 by “photomultiplier” tube.

74. Large Scale Lab Scintillator
The most versatile method for measuring radiation in the laboratory is the scintillation counter.
In this device, a substance called a phosphor emits a flash of light when struck by radiation.
The number of flashes are counted electronically and converted into an electrical signal. 74

75. Scintillator we will use in Lab

76. Sources of Natural Radiation
Radon
A natural decay product of Uranium
Cosmic radiation
From the sun and outside the solar system
Altitude increases exposure, less air protection
Similar to X-Rays
Food
Carbon-14, Potassium-40 are naturally radioactive
A few other elements in trace amounts (e.g. Polonium)
From Earth itself
Uranium, thorium, other natural radioactive elements

78. Radon Alpha emitter, the most ionizing radiation
Short range but very dangerous internally
Radon decay products also radioactive
Solids, not gases …stay on lung tissues

79. Deaths per year, including Radon http://elitehomeinspections.net/radon

80. http://discoveringsomethingneweveryday. blogspot

81. How Radon enters your home

83. Radon exposure in USA http://www.tjernlund.com/radonvac.htm

84. Worldwide Radon distribution http://www. fixradon

86. Source of Radon Danger http://www.grundyhealth.org/newsite/radon.shtml

87. Smoking is greatest cancer risk, Radon is #2

88. Internal & External Radiation
Eating food involves internal exposure
C-14, K-40 in food, cannot be avoided
Cosmic rays from space hard to avoid
Greater risk at higher altitudes, aircraft rides
Ground radiation, radioactive elements
Uranium, Thorium, other natural materials
Radiation is everywhere!
From above, below, and in our food & water

89. X-rays from below the Earth

90. Radiation Units of Measure

91. Differences between radiation units

93. Activity versus dosage not all radiation absorbed … different units used

94. Today’s Experiment We will have 3 stations for 3 tests
(1) Effect of shielding
(2) Isotope half-life
(3) Radiation from household objects
3 or 4 groups of students circulate
Run each test, one after another (any order)
Allocate 20 minutes per test
Pre and post lab questions also included

95. Today’s Experiment, part 1
Effectiveness of Shields
Different radiation has differing penetration
Alpha is comparatively massive, low penetration
Beta is low mass, modest penetration
Gamma has zero mass, high penetration
Intensity reduced by mass of shield
We will use paper, plastic, aluminum, lead,etc.
Distinguish radiation by influence of shields

96. Today’s Experiment, part 2
Half life of an isotope
Intensity of radiation decreases with time
We prepare an isotope with ½ life of few min.
Measure radiation over period of 7 minutes
Plot radiation versus time
We will find half life experimentally

97. Today’s Experiment, part 3
Radiation from common objects
Natural and manufactured items
Radiation sources allaround us
Rocks, minerals, natural objects
Pottery, Smoke detectors
Power plant disasters (Russia, Japan, USA)
Nuclear testing fallout
We will measure several objects
Compare radiation types and amounts

98. Your personal exposure profile
Last part of experiment, done at home
Your assessment of radiation exposure
Medical x-rays (scans, dental, bone density)
Where you live influences exposure
California better than most other places
How often to you use airlines

99. To the Benches! Choose groups of 3-4 Run one station (any order)
Move to next station
Last station … and you’re done.

100. Los Alamos National Laboratory's Periodic Table

101. Marie & Pierre Curie, 1905 Separated tons of mineral pitchblende to discover and isolate Radium, and polonium (named for Poland). The standard measure for radioactivity is the Curie = Ci
101

102. 14C is continually produced in the atmosphere and incorporated into life cycles, so 14C amount in living things is constant. Upon death, no more 14C is absorbed, so concentration decreases. Measuring the remaining radioactivity provides an age estimate. The half life of 14C is 5730 years. The method is good for estimating age of objects 500 to 50,000 years old.
102

103. 14C is continually produced in the atmosphere and incorporated into life cycles, so 14C amount in living things is constant. Upon death, no more 14C is absorbed, so concentration decreases. Measuring the remaining radioactivity provides an age estimate. The half life of 14C is 5730 years. The method is good for estimating age ofobjects 500 to 50,000 years old.
103

104. Radiation exposure http://enhs. umn

105. Radon radiation density in USA http://en. wikipedia

106. Beta Particle Unstable isotope emits electron having negative charge, nucleus of parent changes charge in the other direction, creating a new element. Example is Carbon-14 used for dating historical objects. Note another particle with zero mass and charge, the “Neutrino” … more about that later
106

107. Positron Emission Unstable nuclei can also emit a positive electron, a form of “anti-matter” which turns into energy upon meeting it’s opposite. Mass and charge must balance
107

108. Today’s Experiment, part 1
Radiation versus Distance
Intensity falls off as square of distance
Assumes a “point source” (e.g. sun in sky)
Our source is not exactly a point
Falloff will be approximately linear
You will plot radiation versus distance
A non-linear relationship
Replot with inverse distance
Approximately linear relationship

109. Radiation units of measure Original definition of “curie” was radiation from one gram of radium

114. Emission of g rays causes no change in mass or atomic number.
g emission usually accompanies emission of other rays but it is often omitted from nuclear equations.
Their penetrating power makes them both dangerous to humans and useful in medical applications.
114

115. Proton Emission loss of positive charge reduces atomic number by 1
115

116. Neutron Emission Element stays the same, but new isotope weighs less
116

117. Atomic Bomb Explosion Automatic Camera situated 7 miles from blast with 10 foot lens. Shutter speed 1/100,000,000 second. Joshua tree’s near base vaporized in microseconds.
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