# Chapter 3 - Atomic Mass Units A way to assign relative...

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Atomic Mass Units A way to assign relative atomic masses to elements (much easier than dealing with absolute atomic masses). The 12 C isotope was chosen as a reference point, and assigned a value of 12.0000 amu (atomic mass unit, or “daltons”). So, 1amu = 1/12 the mass of 12 C. All other elements are assigned amu values relative to 12 C. Examples: 1 H has an amu of 1.008 amu 14 N has an amu of 14.007 amu (Why aren’t these amu values exact whole numbers? Because of nuclear binding energy “mass effect” of ≈ 0.01amu) 1amu = 1.661x10 -24 g, which is ≈ the mass of a proton or neutron. Therefore, the approximate mass of an atom is the mass number A . This is not exact, however, due to fractional isotope abundances, the average masses of which give us Atomic Mass/Weight ……

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Natural lithium is: 7.42% 6 Li (6.015 amu) 92.58% 7 Li (7.016 amu) (7.42 x 6.015) + (92.58 x 7.016) 100 = 6.941 amu Average atomic mass of lithium:
Average atomic mass (6.941)

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Atomic Mass/Weight = the average * of the masses of an element’s naturally-occurring isotopes weighted according to their isotopic abundances. Example Problem 1 (solve on board): Find the atomic mass/wt of silver (Ag), given the following data: 51.84% 107 Ag (106.90509 amu) 48.16% 109 Ag (108.90476 amu) Example Problem 2 (solve on board): Boron (Z=5) has two naturally-occurring isotopes. Calculate the %-abundances of 10 B and 11 B, given the following data: Atomic mass of B = 10.81 amu Isotopic mass of 10 B = 10.0129 amu Isotopic mass of 11 B = 11.0093 amu (hint: we are trying to derive two unknowns!) * This is why atomic masses are often close to, but not exactly, 2x the atomic number of an element (because of varying number of isotopes, in varying abundances, in a sample of an element): Ex: C has Z=6 and A=12 (12 amu) U has Z=92 and A=238 (238 amu)
Avogadro’s Number – The Mole Need a way to tangibly connect the macroscopic world to the particulate world of atoms, molecules, ions, etc…. Same reference point as used in amu: 12 C Introducing…..the mole : the amount of a substance that contains as many particles/entities (atoms, molecules, formula units, ions, electrons, etc.) as there are 12 C atoms in exactly 12g of the 12 C isotope. This number of particles/entities is called Avogadro’s Number :

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The mole (mol) = N A = 6.022 x 10 23 (Avogadro’s Number) The units of Avogadro’s Number are mol -1 or entities/mole
1 mol contains 6.022x10 23 particles/entities: 1 mol of 12 C contains 6.022x10 23 12 C atoms 1 mol of H 2 O contains 6.022x10 23 H 2 O molecules 1 mol of NaCl contains 6.022x10 23 NaCl formula units 1 mol of electrons contains 6.022x10 23 electrons 1 mol of students contains 6.022x10 23 students…. A mole is not merely the number 6.022x10 23 . A mole of something contains 6.022x10 23 units/entities of that something. A mole represents a number and a mass …….

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Molar Mass Depending on the atomic mass of an element, the molar mass (the mass of 6.022x10 23 atoms of that element) will be different for each element: 1 Fe atom has an average atomic mass of 55.85 amu 1 mole of Fe atoms has a molar mass of 55.85 g/mol 1 Na atom has an average atomic mass of 22.99 amu 1 mole of Na atoms has a molar mass of 22.99 g/mol 1 Cl atom has an average atomic mass of 35.45 amu 1 mole of Cl atoms has a molar mass of 35.45 g/mol So, for any element, the numerical value of the atomic mass (in amu) will equal the numerical value of the molar mass (in g/mol).
One Mole of: C S Cu Fe Hg

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1 amu = 1.66 x 10
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