1. Unit 12 Mass and Moles

2. Avogadro’s Hypothesis
Under the same conditions of temperature and pressure, equal volumes of two different gases will have the same number of particles regardless of mass.
At STP, this number of particles is Avogadro’s number or 6.02 x 1023 or 1 mole

3. *YOU MUST MEMORIZE THIS*
This number can be used to express the mass, volume or number of particles for any gas sample. The conversion from one to another involves this equivalency:
1mole = 6.02 x 1023 particles = gram formula mass = 22.4 liters of space
*YOU MUST MEMORIZE THIS*

4. Determining mass using formulas
Atomic Mass Unit (amu)
1 amu = 1/12 the mass of a C-12 atom
Gram Atomic Mass
The mass of 1 mole of atoms
Numerically equal to atomic mass number
The unit used is (g) gram

5. Gram Molecular Mass
The mass of 1 mole of molecules
Found by adding up the atomic mass numbers of each element in the molecule
The unit used is (g) grams

6. Gram Formula Mass
The mass of 1 mole of an ionic substance
*remember that ionic substances are not composed of molecules!
Determine the gram formula mass/gram molecular mass of the following:

7. NaCl
H2O
1 mole of H2 gas

8. Avogadro’s Number and Mole Map Problems
What we know:
At STP and in the gaseous phase:
1 mole = 6.02 x 1023 particles of a substance
1 mole takes up 22.4 L of space
1 mole has a mass equal to the gram formula/gram molecular mass

9. We could convert from moles to mass to volume using a Mole Map to set up a proportion
Example: What is the mass of 3.01 x 1023 molecules of NH3(g) at STP?

10. What is the mass of 5 moles of O2 gas at STP?

11. How many molecules of NH3 would take up 44.8 L at STP?

12. What is the mass of 9.03 x 1023 atoms of neon gas at STP?

13. Density The mass per unit volume of a substance.
The density of a substance DOES NOT change for that substance!
Formula: Density = Mass/Volume

14. 3 ways for the Regents to ask Density problems
Table S: What is the density of one mole of N2?
 Question will give 2 variables, you must use the formula to solve for the 3rd_: What is the density of a solid having a mass of 75g and a volume of 3 cm3?

15. Given the density, the Regents can ask for the molecular mass at STP: Remember that at STP, the molecular mass or 1 mole of a gas has what volume? 22.4L
What is the gram molecular mass of a gas having a density of 2g/L at STP?

16. Percent Composition
To determine the percentage by mass of a particular element in a compound.
Determine the gfm
Take the mass of individual element and divide by gfm
Multiply by 100%

17. Examples What percent by mass of CaCO3 is made up of calcium?
What is the percent by mass of nitrogen in NH4NO3
What is the percent by mass of water in copper sulfate pentahydrate?

18. Empirical Formula from Percent Composition
Using the percent composition, one can determine the smallest whole number ratio of atom to atom in a compound (empirical formula). **Remember that since ionic substances do not have true molecules, they are always expressed in empirical formulas**
Follow the steps:
Divide the percent by the atomic mass
Take the answer to step 1 and divide it by the smallest answer to step 1.

19. Examples
What is the empirical formula of a compound containing 40% calcium, 12% carbon and 48% oxygen by mass?
What compound contains 56.58% potassium, 8.68% carbon and 34.73% oxygen by mass?

20. Special Rule
if the answer to step 2 ends in .1 or .9 you can round off, BUT if it ends in .3 or .5, this is a significant portion of a number. You cannot round away this number, instead you must adjust all numbers accordingly by following the formula:
If the number ends in .3 _multiply everything by 3____
If the number ends in .5 __multiply everything by 2___

21. Example
Give the empirical formula of a compound containing 90.7% lead and 9.3% oxygen by mass.

22. Molecular Formula from Empirical Formula
To determine the molecular formula from the empirical formula follow the steps: *Calculate the empirical mass. *Divide the mass of the compound by the mass of the empirical formula. *Multiply all subscripts by the answer in step 2.

23. Example
A compound has a molecular mass of 42 amu and an empirical formula of CH2. What is the molecular formula?

24. Moles in Balanced Equations
A chemical equation usually represents a chemical reaction. The equation will identify:
The reactants and products
The molar ratio of each of these.
Phases of matter for each substance.
Possibly some reference to energy changes in the reaction.

25. 2H2(g) + O2(g) → 2H2O(l) + heat
Example
2H2(g) + O2(g) → 2H2O(l) + heat

26. When using equations to solve conversion problems, one must remember the proportions the original substances are in. One can determine how many grams of a reactant are needed to produce a set volume by using the following rules:

27. Mass is the ONLY part of Avogadro’s hypothesis that CAN NOT be used directly in a proportion!!
You must convert grams to moles–Table T
Follow these steps:
read and underline
cross out what is not involved
set up a proportion coefficients/known information
cross multiply and divide

28. Example
What volume of CO2(g) is produced when 15 liters of O2(g) are consumed in the reaction:
C2H4(g) O2(g) → CO2(g) H2O(g)

29. Given the reaction:
2 C2H6(g) O2(g) → CO2(g) H2O(g)
What number of carbon dioxide molecules are produced when 6.02x1023 molecules of ethane are consumed?