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| Elements are composed of small paricles called atoms. |
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| All atoms of one element are identical. Atoms of one element differ from all other elements. |
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| two or more element atoms combine in fixed whole number ratios to form chemical compounds. |
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| Atoms are not created or destroyed in regular chemical reactions. |
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| Law of conservation of mass |
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| Lavoisier, 1743-94, Matter cannot be created or destroyed in regular chemical reactions. |
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| Law of Definite Proportions |
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| Proust 1754-1826, A chemical compound always contains the same proportion of elements by mass. A compound has the same composition regardless of origin. |
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| By shooting alpha particles at gold foil rutherford was able to see that the atomic model is comprised of a very small, very dense, positivly charged nucleus. Electrons comprise most of the space of the atom but very little mass. Atoms are mostly empty space. The number of protons is = to the # of electons. The nucleus contains protons = to its atomic number and neutrons account for the rest of the mass of the atom. |
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| Ionized oil drop experiment where negatively chaged oil droplets where injected into a space between charged disks. Determined the charge of the electron. |
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| JJ Thomson experiment/ model |
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| Used cathode ray tubes, electromagnets, and electrodes to find that cathode rays are comprised of electrons, that electrons are the fundamental negative subatomic particle, the charge to mass ratio of elctrons. (z/g) is the same regardless of the source of rays. Plum pudding model disproved by rutherford. |
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| the atomic model is comprised of a very small, very dense, positivly charged nucleus. Electrons comprise most of the space of the atom but very little mass. Atoms are mostly empty space. The number of protons is = to the # of electons. The nucleus contains protons = to its atomic number and neutrons account for the rest of the mass of the atom. |
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| given isotopic mass and fractional abundance, calculate atomic weight (amu) |
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| isotopic mass X percent abundance in decimal form. Add these numbers to find average mass of all isotopes. |
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| given mass, calculate moles. visa versa |
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mass X moles/atomic weight= #of moles
moles X atomic weight/moles = mass.
Atomic mass = 1 mole = 6.022 X 10 to the 23 atoms |
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| charged atom made because an atom or chemically bonded group of atoms lost or gained electrons |
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| atoms of an element that have different mass due to a different # of electrons. |
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| 6.022 X 10^23 number of carbon 12 atoms in exactly 12 grams of carbon 12 |
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| fraction of the total # of atoms of an element that are a certain isotope |
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| the weighted average of all the naturally occuring isotopes of the element (relative to carbon 12 |
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| the amount of a substance that contains the same number of elementary units as there are carbon 12 atoms in exactly 12 grams of carbon 12 |
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| the mass in grams of one mole. The molar mass for any element is = to the atomic weight |
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