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The Mole _Avogadro's number - slider-dpchemistry-11

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									Y11 DP Chemistry
    R. Slider
As you know, atoms are very, very small…

            Some single atom masses:
            Li: 1.15217 x 10-23 g    O: 2.65659 x 10-23 g
            He: 6.64605 x 10-24 g    U: 3.95233 x 10-22 g


These masses are not terribly convenient to work with, so chemists work with
relative masses. The relative atomic masses of all the elements are based on the
mass of C-12 (1.99X10-22 g), the most abundant isotope on the Earth. In fact:

          Relative atomic mass (Ar) is defined as the mass of an
          element relative to 1/12 of the mass of an atom of the
          Carbon-12 isotope.


  These relative atomic masses are conveniently located on the Periodic Table
  and have no units since they are all relative. H: 1.01 P: 30.97 Br: 79.91
The mole is a unit used only in
Chemistry that denotes the amount
of substance (n).

Since chemists had determined the
relative atomic masses of the elements, it
logically followed that 1.01g of H had
the same number of atoms as 12.0g of C.
And that same number of particles is                  This is also a mole…
equal to a mole of that substance.

The Mole is…
The number equal to the number of carbon atoms in exactly 12 grams of
pure 12C. (We once compared to H, the lightest element. Now we use C-12,
the most abundant)
 1 mole of anything = 6.022  1023 units of that thing (atoms, ions,
molecules, grains of sand, etc.)
  12g of C-12 contains 6.022x1023 atoms.
      This is known as Avogadro’s
      Number (NA or L) and is equivalent
      to 1 mole of carbon atoms.                 Substance     Molar mass
                                                               (g/mol)
Avogadro’s number = 6.02x1023 = 1mole           carbon              12
                                                chlorine           35.5
 Notice that the mass of 1 mole of C-12 is
 the same value as the relative atomic mass     oxygen              16
 for C-12 on the Periodic Table.
                                                water               18
 In the same way, 1 mole of any element or
 compound is equivalent to its atomic,
 molecular or formula weight.
 We can now define the relative atomic        Molar mass = mass of 1 mole of
 masses from the periodic table as molar                  any substance
 masses with the units g/mol.
From the previous example using C-12, we now
  have a mathematical relationship between
  mass and moles, which is:

                         Mass of substance (g)
  Number of moles (n) =
                        Molar mass (MM) (g/mol)


Example: How many moles are in 25 g of CO?
          n CO = 25g CO/(12+16)g/mol
                = 0.89 moles
   We can also convert between moles and the
    number of atoms or molecules using
    Avogadro’s number
       Number of atoms/molecules = moles (n) x NA


Example: How many atoms are there in a copper pipe that
  weighs 2.56g?
              n = 2.56/63.6 = 0.0403 moles
       number of Cu atoms = 0.0403 X 6.022X1023
                      = 2.43X1022
  We can now use the mole concept to determine
   how much product to expect in a chemical
   reaction. Take the following example:

             2Fe2O3(s) + 3C(s)  4Fe(l) + 3CO2(g)

The coefficients in front of each species provide us with useful ratios that
  we can use to calculate expected masses of products in a chemical
  reaction. We previously said that these were ratios based on the
  numbers of atoms. However, with Avogadro’s number, we can now say
  that these are molar ratios.
We say, 2 moles of iron (III) oxide react with 3 moles of carbon to produce
  4 moles of iron and 3 moles of carbon dioxide gas.
               2Fe2O3(s) + 3C(s)  4Fe(l) + 3CO2(g)
Example:
How many grams of iron will we expect if we react 12g of
  iron (III) oxide as in the above reaction, assuming
  neither reactant is in excess?
                                                   Convert to g of
        Convert to moles         Molar ratio          unknown



12g Fe2O3 X   1 mol Fe2O3        4 mol Fe           55.9g Fe
                             X                 X
              159.8g Fe2O3       2 mol Fe2O3        1 mol Fe

      12 X 1 X 4 X 55.9
 =                           =    8.4 g Fe
      159.8 X 2 X 1
 Molecular mass is the sum of the atomic masses
  of the atoms in a molecular formula.


Example:
The molecular mass of sucrose (table sugar) C12H22011 is
  calculated as:
M.W. = (12XAC) + (22XAH) + (11XAO)
M.W. = (12X12.0) + (22X1.01) + (11X16.0)
     = 342.2
Formula mass is the sum of the atomic masses in a
   compound which has no discreet molecules (e.g. ionic
   compounds). These describe the ratios of the atoms
   present (i.e. empirical formulas), but are calculated
   the same way as molecular masses.


Example:
The formula mass of calcium phosphate Ca3(PO4)2 is
  calculated as:
F.W. = (3XACa) + (2XAP) + (8XAO)
F.W. = (3X40.1) + (2X31.0) + (8X16.0)
     = 310.3
1.   What is the symbol for relative atomic mass?
2.   Write the relative atomic masses for O, Mg, S.
3.   Calculate the molecular mass of ethanol C2H5OH.
4.   Calculate the molar mass of copper(II) sulfate CuSO4.
5.   State the value and symbols of Avogadro’s Number.
6.   How many atoms are there in 3 moles of nitrogen
     atoms? 3 moles of ammonia?
7.   How many grams are there in 3 moles of nitrogen? 3
     moles of ammonia?
8.   How many moles are there in 48g of water?

								
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