Nuclear Chemistry, Atoms: The Building Blocks of Matter

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Transmutation
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The change in the identity of an isotope due to a change in the number of its protons
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Alpha Decay definition
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When an isotope loses a particle with a 2 proton and a mass of 4
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Beta Decay definition
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When an isotope gains a proton through the loss of an electric charge on one of its neutrons
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Gamma Decay definition
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When an isotope decays to release only energy, not a particle
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Positron
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a positive particle with the same mass as an electron given off as a result of a proton changing to a neutron
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Half-life
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the amount of time it takes for half of a substance to decay into another substance
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Electron
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a negative sub atomic particle
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Fission
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the splitting of one large nucleus to create two smaller nuclei (two new elements with some mass being converted into energy)
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Fusion
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the combining of two smaller nuclei to create a larger nucleus (some mass is converted to energy)
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Carbon Dating
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the process using Carbon-14 to date materials that were once alive
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Ion
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a charged atom
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isotope
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an atom with fewer or more neutrons than the average form of that element
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Alpha decay example
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Beta decay example
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Gamma Decay example
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Atom
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“indivisible” (Greek meaning), the smallest particle of an element that retains the chemical properties of that element
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Dalton’s Atomic Postulate 1
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All matter is made of atoms and atoms are indivisible, cannot be subdivided, created or destroyed
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Dalton’s Atomic Postulate 2
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All atoms of the same elements are identical in mass and properties (the same)
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Dalton’s Atomic Postulate 3
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Compounds are combinations in whole number ratios of two or more types of atoms
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Dalton’s Atomic Postulate 4
Dalton's Atomic Postulate 4
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In chemical reactions, atoms are combined, separated or rearranged
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Dalton’s Atomic Postulate 5
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Compounds are made in definite proportions
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Law of Conservation of Mass
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(Law of Lavoisier) in chemical reactions, the total mass is conserved
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Law of Definite Proportion
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(Proust) when a chemical compound is formed, their is a definite proportion of the atoms forming the compound (example: in table salt; sodium chloride; it ALWAYS has by mass, the same amount of sodium (Na , 39.34%) and chlorine (Cl, 60.66%)
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Law of Multiple Proportions
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if 2 or more different compounds are composed of the same 2 elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ration of small whole numbers. (carbon and oxygen form carbon dioxide and carbon monoxide; the ratio of carbon to oxygen between dioxide to monoxide is always 2:1)
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Atomic Number
Atomic Number
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number of protons of each atom of that element
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Isotope
Isotope
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atoms of the same element that have different numbers of neutrons in the nucleus (so different atomic masses) Tin has the most isotopes (10); although they differ in masses isotopes do not differ significantly in their chemical behavior
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Mass number
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the total number of protons plus neutrons that make up nucleus of an isotope
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Nuclide
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general term for a specific isotope of an element
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Atomic Mass Unit
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(amu) a unit of mass that describes the mass of an atom or molecule
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Mole
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(mol) the SI unit for amount of substance: the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12
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Avogadro’s number
Avogadro's number
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the number of particles in exactly one mole of a pure substance (6.022 1415 x 10^23 or rounded to 6.022 x 10^23 = exactly what 12g of carbon-12 atoms contains)
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Molar mass
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equal to the atomic mass of the element; the mass of one mole of a pure substance is called the molar mass of that substance; written in units of g/mol; FOUND ON THE PERIODIC TABLE FOR EACH ELEMENT
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Subatomic particles
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what make up an atom; the three main are protons,neutrons, and electrons
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Chemical Reaction
Chemical Reaction
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transformation of a substance or substances into one or more new substances
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Nucleus
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a very small region located at the center of an atom
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Proton
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positively charged (+) particle equal in magnitude to the negative charge of an electron; at least one in every atom
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Neutron
Neutron
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neutral particle; at least one or two in every atom
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Electron
Electron
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negative charged particle (-); surrounds the nucleus in an occupied region
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Cathode-ray tubes
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glass tubes; in late 1800s, experiments were performed where electric current was passed through various gases at low temp; the resulting glowing stream was called cathode rays
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Average Atomic Mass
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weighted average of the atomic masses of the naturally occurring isotopes of an element
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JJ Thomson
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(1897) discovered electrons, charge to mass ratio and that all cathode rays are composed of identical negatively charged particles named electrons
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Rutherford
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(1911) discovered the nucleus, protons; Gold foil experiment
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Chadwick.
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(1934) discovered the neutrons
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Nuclear forces
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the short-range proton-neutron, proton-proton, and neutron-neutron forces hold the nuclear particles together
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Neutron equation (find # of neutrons)
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Mass # – atomic # = # neutrons (protons+neutrons)-protons=neutrons
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cathode rays
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streams of negatively charged particles (electrons).
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plum pudding model (Thomson)
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JJ Thomson; negative electrons are spread evenly through the positive charge of the rest of the atom (think watermelon)
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Robert A. Millikan (1909)
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1. since atoms are electronically neutral they must contain a positive charge to balance the negative electrons 2. because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass
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1911 – Rutherford, Geiger and Marsden
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bombarded a thin piece of gold foil with a narrow beam of alpha particle, some particles were deflected back to source – reasoned this was the nucleus and it was a very small part of atom because the rest of the particles passed through undisturbed
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Nuclei
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different elements differ in number of protons (positive charge) so the number of protons determines the atoms identity
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nuclear forces
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the short range proton-neutron; proton-proton and neutron-neutron forces that hold the nuclear particles together
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picometer (pm)
picometer (pm)
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Radius of an atom from the center of the nucleus to the outer portion of the electron cloud ; 1 pm = 10^-12(ten to the negative 12th power) meters or 10^-10 centimeters
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atomic number
atomic number
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number of protons in each atom of that element
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Designating isotopes
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Identified by specifying their mass number in 2 ways: 1. Hyphen notation: hydrogen – 3 2. Nuclear symbol: 3 H 1 with mass number (neutron + protons) on top and atomic number (protons) on bottom
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atomic mass units (amu)
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all atoms compared to carbon-12 atoms (12 amu) so 1 amu = 1/12 of carbon
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calculating average atomic mass
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multiply mass of each type by the decimal fraction representing its % in the mixture 25% = .25 each weighing 2g; 75%=.75 each weighing 3g so (2g*.25)+(3gx.75)=2.75g
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gram/mole conversions
gram/mole conversions
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(moles of substance) x molar mass = grams and (moles)(gram/moles)=grams
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Rutherford model
Rutherford model
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1911 model of the atom. small nucleus surrounded by electrons in orbit around it
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Bohr atomic model
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Suggest that electrons move in a definite path around the nucleus. 1913
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Quantum mechanical model (electron cloud model)
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a mathematical model and most accurate model of an atom used today
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Dalton model
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The first major model of the atom, developed in 1800.
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Greek model
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Basically, this model was a philosophical idea that everything could be broken down into one fundamental unit — the atom. It could not be tested at that point in time because the Greeks did not posses the technology to observe anything that was so microscopic such as the atom.
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nucleons
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the protons and neutrons in atomic nuclei
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nuclide
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the reference to an atom in nuclear chemistry; identified by the number of protons and neutrons in its nucleus
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mass defect
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difference between the mass of an atom and the sum of the masses of its protons, neutrons and electrons
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what causes mass defect?
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conversion of mass to energy upon formation of the nucleus
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nuclear binding energy
nuclear binding energy
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the energy release when a nucleus is formed from nucleons; E=mc^2
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binding energy per nucleon
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is the binding energy per of the nucleus divided by the number of nucleons it contains – the higher the binding energy the more tightly the nucleons are held together and are therefore more stable
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band of stability
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stable nuclei cluster over a range of neutron-proton ratios (the graph of elements and their isotopes plotted)
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nuclear shell model
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nucleons exist in different energy levels or shells in the nucleus
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magic numbers
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Numbers of nucleons that represent completed nuclear energy levels: 2, 8, 20, 28, 50, 82, and 126 – the most stable nuclides
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nuclear reaction
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reaction that affects the nucleus of an atom
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transmutation
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change in the identity of a nucleus as a result of a change in the number of its protons
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radioactive decay
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the spontaneous disintegration of a nucleus into a slightly lighter nucleus, accompanied by emission of particles, electromagnetic radiation or both
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Henri Becquerel
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1896 wrapped a photographic plate in lightproof covering and place uranium on top of it. Figured out didn’t need sunlight to expose the plate, it was the radioactive decay of uranium that exposed the plate
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nuclear radiation
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particles or electromagnetic radiation emitted from the nucleus during radioactive decay
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radioactive nuclide
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an unstable nucleus that undergoes radioactive decay – all nuclides beyond atomic #83 are unstable and radioactive
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alpha particle
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two protons and two neutrons bound together and emitted from the nucleus during some kinds of radioactive decay; charge is 2+, represented by symbol
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beta particle
beta particle
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an electron emitted from the nucleus during some kinds of ratioactive decay. atomic number 1+, mass stays the same
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positron
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particle that has the same mass as an electron but has a postive charege, and is emitted from the nucleus during some kinds of ratio active decay
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electron capture
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an inner orbital electron is captured by the nucleus of its own atom. the inner orbital electron combines with a proton and a neutron is formed; atomic number decreases by one but mass number doesn’t change (stays the same)
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gamma rays
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high energy electromagnetic waves emitted from a nucleus as it changes from an excited state to a ground energy state
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penetration of particles
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alpha: your hand beta: aluminum gamma: lead neutrons: concrete
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half life
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time required for half of the atoms of a radioactive nuclide to decay
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decay series
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a series of ratioactive nuclides produced by successive radioactive decay until a stable nuclide is reached
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parent nuclide
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heaviest nuclide of each decay series
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daughter nuclides
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nuclides produced by the decay of the parent nuclides
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artificial radioactive nuclides
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radioactive nuclides not found naturally on earth
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artificial transmutations
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bombardment of nuclei with charged and uncharged particles to make artificial radioactive nuclides; radioactive isotopes of all the natural elements have been produced this way
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transuranium elements
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elements with more than 92 protons in their nuclei – all are radioactive
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roentgen (R)
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unit used to measure nuclear radiation exposure
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rem
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unit used to measure the dose of any type of ionizing radiation that factors in the effect that the radiation has on human tissue
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film badges
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use exposure of film to measure the approximate radiation exposure of people working with radiation
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Geiger-Muller counters
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are instruments that detect radiation by counting electric pulses carried by gas ionized by radiation; typically used to detect beta particles, x rays and gamma radiation
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scintillation counters
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instruments that convert scintillating light to an electrical signal for detecting radiation
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radioactive dating
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process by which the approximate age of an object is determined based on the amount of certain radioactive nuclides present: Carbon 14 is radioactive has half life of approximately 5715 years can be used to estimate the age of organic material to about 50,000 years
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radioactive tracers
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radioactive atoms that are incorporated into substances so that movement of the substances can be followed by radiation detectors – used to diagnose cancer etc
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nuclear fission
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nucleus of very heavy atom (uranium) split into 2 or more lighter nuclei; releases high amounts of energy
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chain reaction
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reaction in which the material that starts the reaction is also one of the products and can start another reaction
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critical mass
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minimum amount of nuclide that provides the number of neutrons needed to sustain a chain reaction
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nuclear reactor
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use controlled fission chain reactions to produce energy and radioactive nuclides
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nuclear power plants
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use energy as heat from nuclear reactors to produce electrical energy
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shielding
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radiation absorbing material that is used to decrease exposure to radiation especially gamma rays from nuclear reactors
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control rods
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neutron absorbing rods that help control the reaction by limiting the number of free neutrons
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moderator
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used to slow down the fast neutrons produced by fission
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nuclear fusion
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low mass nuclei combine to form a heavier more stable nucleus; opposite of fission; creates even more energy; cannot currently be controlled due to heat; named and explained by Lise Meitner

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