1. Unit 5 ~ The Mole (Chapter 8)
And you

2. 5-1 The Mole (Sections 8.1 – 8.3)  
The atomic mass on the period table is the weighted average of all the elements’ isotopes masses. The unit for a single atom of an element is the amu or atomic mass unit. The amu is an incredibly small number of grams, since atoms are so small:
1 amu = 1.66 x g.

3. Obviously we can’t weigh out a single atom
Obviously we can’t weigh out a single atom. However, we can weigh a large number of atoms and use this mass to determine the number of atoms present. It’s analogous to purchasing a bag of small food, like beans or rice, by weighing one pound of the food; if you know the mass one bean, you know how many beans you’ve purchased. Another example, let’s say a penny has a mass of 2.5 g, and you want to count a huge pile of pennies, you could simply find the mass and multiply by (1 penny/2.5 g).

5. To relate the mass in grams of a single atom, take the inverse of 1
To relate the mass in grams of a single atom, take the inverse of 1.66 x The answer = x This means that if you take x 1023 atoms, the value of is scaled up to 1 gram. This number, which has been defined as the number of atoms in g of Carbon, is referred to as “the Mole” or Avogadro’s Number. It allows chemists to use the periodic table to count atoms by weighing, just like the pennies above.

6. Important Concepts: 1 atom of Carbon = 12.011 amu
1 mole of Carbon = grams = x 1023 atoms of Carbon
1 atom of Mg = amu
1 mole of Mg = grams = x 1023 atoms of Mg
The mole is simply a VERY large number, used to count atoms and molecules by weighing.
It’s analogous to “1 dozen = 12”. The mole = x 1023, and can be used for anything.

8. 5-2 The Mole Bridge (Section 8.3)
The mole “bridge” is a schematic that is useful in converting between the mass of a substance, the number of moles of the substance, and the # of atoms or molecules of the substance:

9. 1 mole = mass on PT 1 mole = x 1023
To use the mole bridge, simply start with what you’re given, and move either 1 or 2 bridges to where you want to be. Each “bridge” represents a conversion factor that can be inverted to cancel units using dimensional analysis.
Mass (g)
Moles
Atoms or Molecules

10. Example 1: 1 bridge 33.6 g Fe = ? Moles
33.6 g Fe x 1 mole Fe = mole Fe
55.85 g

11. Example 2: 1 bridge 1.20 x 1024 molecules water = ? moles water
1.20 x 1024 molec. H2O x 1 mole = x 1023 molec. H2O
1.99 mol H2O

12. Example 3: 2 bridges 2.4 x 1024 atoms Ca = ? g Ca
2.4 x 1024 atoms Ca x mol Ca x g Ca = x 1023 atoms Ca mol Ca
159.7 or 160 g Ca
Now you can see how we can determine how many atoms, and how many moles of atoms we have by measuring the grammage!
No “grammage” is not a real word that you will use outside of this class – so do not use the term in Mr. Albritton’s chem. II class next year – or he will be unhappy

13. Example 4: 2 bridges 110.5 g Pb = ? atoms Pb
You try this one on your own please!!!!
110.5 g Pb x mole Pb x x 1023 atoms Pb =
207.2 g Pb mole Pb
3.212 x 1023 atoms Pb
Well done!!!!!

14. Important notes:
1) Answer has same number of sig figs as original measure.
2) The conversion factors always have “1 mole” (in this type of conversion problem)

15. 5-3 Molar Mass (Section 8.4)
The Molar Mass (M) is the mass in grams of 1 mole of a substance; units are g/mole.
Synonomous terms to Molar Mass are relative Molar Mass (Mr) and the older term Formula Weight (FW).
To calculate the Molar Mass simply add the atomic masses of the elements in a formula:
M =  atomic masses = mass of 1 mole
For this class, let’s report M to 0.01 g.

16. Practice: H2O = 2 mol H + 1 mol O = 2(1.01g) + 16.00 = 18.02 g/mol H2O
That was easy and more than a little fun!!!!

17. Al2(CO3)3 = 234.03 g/mol 2 mol. Al + 3 mol C + 9 mol O = 9(16.00)
2(27.00)
3(12.01) = + +
g/mol

18. CuSO4•5H2O
1 Cu S O (10 H O)
[ (16.00)] + [10(1.01) + 5(16.00)] =
g/mol CuSO4•5H2O (copper II sulfate hydrate)