apchapt3 - Atomic Mass Atoms are so small, it is difficult...

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Atomic Mass Atoms are so small, it is difficult to discuss  how much they weigh in grams. Use atomic mass units. an atomic mass unit (amu) is one twelth the  mass of a carbon-12  atom. This gives us a basis for comparison. The decimal numbers on the table are  atomic masses in amu.
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They are not whole numbers  Because they are based on averages of  atoms and of isotopes. can figure out the average atomic mass  from the mass of the isotopes and their  relative abundance. add up the percent  as decimals times the   masses of the isotopes.
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Examples There are two isotopes of carbon  12 C with a  mass of 12.00000 amu(98.892%), and  13 with a mass of 13.00335 amu (1.108%). There are two isotopes of nitrogen , one with  an atomic mass of 14.0031 amu and one with a  mass of 15.0001 amu. What is the percent  abundance of each? 
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The Mole The mole is a number. A very large number, but still, just a  number. 6.022 x 10 23  of anything is a mole A large dozen. The number of atoms in exactly 12 grams  of carbon-12.
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The Mole Makes the numbers on the table the mass of  the average atom.
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More Stoichiometry
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Molar mass Mass of 1 mole of a substance. Often called molecular weight. To determine the molar mass of an element,  look on the table. To determine the molar mass of a  compound, add up the molar masses of the  elements that make it up.
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Find the molar mass of CH 4   Mg 3 P 2   Ca(NO 3 ) 3   Al 2 (Cr 2 O 7 ) 3   CaSO 4  ∙ 2H 2 O
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Percent Composition Percent of each element a compound is  composed of.
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This note was uploaded on 07/24/2008 for the course CHM 101 taught by Professor Geldart during the Spring '07 term at Rhode Island.

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apchapt3 - Atomic Mass Atoms are so small, it is difficult...

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