2. The Mole

3. The Mole Memorize this number: 1 mol = 6.02 x 1023 of something
A mole is defined as the number of particles in exactly 12g of Carbon-12. This number = 6.02 x 1023 particles.
6.02 x 1023 is called Avagadro’s Number after Amadeo Avagardro who determined the volume of 1 mol of gas in 1811
Memorize this number:
1 mol = 6.02 x 1023 of something

4. The Mass of a Mole
Atomic Mass is the relative mass of single atom of an element in amu
Atomic masses on the periodic table can be used to determine molar masses.
Molar Mass is the mass of 1 mole of something in grams
If the atomic mass of carbon is amu, then the molar mass of carbon is g/mol
IN SHORT: The molar mass of any element is exactly equal to its atomic mass – just change the units

5. Moles of Compounds
In a mole of a chemical compound (molecule or formula unit), there is also one or more moles of each atom.
For example:
In 1 mol of CCl2F2 (Freon), there is:
1 mol C
2 mol Cl
2 mol F
If you have 2 mol of MgCl2, there is:
2 mol Mg
4 mol Cl
(Multiply the coefficient x the subscript!)

6. The Mole: In Short
A mole is 6.02 x 1023 of anything, just like a dozen is 12 of anything.
Atomic mass units (amu) is a way of comparing the sizes of different elements.
atomic mass in amu = molar mass in g

7. Using the Mole You should be able to:
Convert Moles to Particles
Convert Particles to Moles
Convert Moles to Mass
Convert Mass to Moles
Convert Mass to Particles
Convert Particles to Mass
Determine percent composition
Determine empirical formula
Determine molecular formula
We are going to practice this a lot!!!!!

8. Preliminary: Moles to Mass of an Element
Determine the numbers of moles of each element in 1 mole of CaF2.
1 mole of Ca
2 mole of F
Determine the number of grams of each element in 1 mol of CaF2.
1 mol Ca x 40.1 g/mol = 40.1 grams Ca
2 mol F x 18.9 g/mol = 37.8 grams F

9. Preliminary: Moles to Mass of Molecule
Determine the number of grams of each element in 1 mol of CaF2.
1 mol Ca x 40.1 g/mol = 40.1 grams Ca
2 mol F x 18.9 g/mol = 37.8 grams F
What is the molar mass of CaF2?
1 mol Ca + 2 mol F = 1 mole CaF2
40.1 g Ca g F = 77.9 g CaF2

10. Preliminary: Mass of a Compound
To find the mass of a compound, simply add up the molar mass of each of its atoms. For example, calculate the molar mass of CO2.

11. Reminders!! Particles can mean: atoms, molecules, formula units, etc.
Covalent compounds = molecules
Ionic compounds = formula units
Use the factor-label method!!!
Learn how to multiply and divide with exponents on your calculator! (Check with me if you need help).

12. The Flow Chart Mass (g) Moles Part Pieces Atoms Molecules Ions Etc.
÷ Avo’s #
x Molar Mass
Mass (g)
Moles
÷ Molar Mass
x Avo’s #

13. #1 Moles to Particles
Determine the number of formula units in 4.3 moles of NaCl.
Multiply the number of moles by the appropriate conversion factor for Avogadro’s number. (Multiply by Avogadro’s #).

14. Practice: Moles to Particles
Determine the number of atoms in 2.50 mol Zn.
Determine the number of formula units (i.e., particles of an ionic compound) in 3.25 mol AgNO3.
Determine the number of molecules in 11.5 mol H2O.

15. Practice: Moles to Particles (x)
Determine the number of atoms in 2.50 mol Zn.
Determine the number of formula units (i.e., particles of an ionic compound) in 3.25 mol AgNO3.
Determine the number of molecules in 11.5 mol H2O.
NOTE: DON’T FORGET SIGNIFICANT DIGITS!!!
1.51 x 1024
1.96 x 1024
6.92 x 1024

16. #2 Particles to Moles
How many moles are there in 5.64x1025 atoms of Fe?
Multiply the number of atoms by the appropriate conversion factor for Avogadro’s number. (Divide by Avogadro’s #).

17. Practice: Particles to Moles
How many moles contain each of the following?
5.75x1024 atoms of Al
3.75x1024 molecules of CO2
3.58x1023 formula units of ZnCl2
2.50x1020 atoms of Fe
3.4x1017 atoms of Na

18. Practice: Particles to Moles
How many moles contain each of the following?
5.75x1024 atoms of Al
3.75x1024 molecules of Co2
3.58x1023 formula units of ZnCl2
2.50x1020 atoms of Fe
3.4x1017 atoms of Na
9.55 mol
6.23 mol
0.594 mol
4.15x mol
5.64x10-7 mol

19. Practice
WS #1 Moles to Particles

20. #3 Moles to Mass
If you have mol of Cr, how many grams do you have?
Determine the molar mass of Cr from the periodic table: g/mol (same number as the atomic mass)
Multiply the number of moles by the conversion factor. (Multiply by the molar mass).

21. Practice: Moles to Mass (x)
Convert from moles to mass:
3.57 mol Al
42.6 mol Si
3.45 mol Co
2.45 mol Zn
8.2 mol Pb

22. Practice Convert from moles to mass: 3.57 mol Al 42.6 mol Si
3.45 mol Co
2.45 mol Zn
8.2 mol Pb
96.3 g
1.2 x 103 g
203 g
1.60 x 102 g
1.7 x 103 g

23. #4 Mass to Moles How many moles of Ca are in 525g of Ca?
Determine the molar mass of Ca from the periodic table: g/mol
Multiply the given mass (525g) by the conversion factor. (Divide by the molar mass).

24. Practice: Mass to Moles
Convert from mass to moles:
25.5 g Ag
300.0 g S
125 g Zn
99g O
1.0 kg Fe

25. Practice: Mass to Moles
Convert from mass to moles:
25.5 g Ag
300.0 g S
125 g Zn
99g O
1.0 kg Fe
0.236 mol
9.355 mol
1.91 mol
6.2 mol
17.9 mol

26. Practice
WS #2 Mole to Mass
WS #3 Mixed One-Step Mole WS

27. Multi-Step Problems How many molecules are there in 24 g of FeF3?
How many grams are there in 7.5 x 1023 molecules of AgNO3?
Calculate the molar mass of FeF3
Divide the mass you have by the molar mass. This gives you the number of moles.
Multiply the number of moles by Avogadro's number.
Calculate the molar mass of AgNO3.
Divide the number of molecules by Avogadro’s number. This gives you the number of moles.
Multiply the number of moles by the molar mass.

28. #5 Mass to Particles
Calculate the number of atoms in a 25.0 g nugget of pure gold (Au).
Convert mass to moles (mass ÷ molar mass)
Convert moles to atoms (moles x 6.02 x 1023)

29. Practice: Mass to Particles
How many atoms are there in 14 g of C?

30. #6 Particles to Mass What is the mass of 5.50 x 1022 He atoms?
Convert atoms to moles (atoms ÷ 6.02 x 1023)
Convert moles to mass (moles x molar mass)

31. Practice: Particles to Mass
What is the mass of 7.6 x 1024 atoms of Ca?

32. Practice: Mass to Atoms and Atoms to Mass
Convert to Atoms
Convert to Mass
55.2 g Li
0.230 g Pb
11.5 g Hg
45.6 g Si
0.120 kg Ti
6.02 x 1024 atoms Bi
1.00 x 1024 atoms Mn
3.40 x 1022 atoms He
1.50 x 1015 atoms N
1.50 x 1015 atoms U

33. Practice: Mass to Atoms and Atoms to Mass
Convert to Atoms
Convert to Mass
55.2 g Li
0.230 g Pb
11.5 g Hg
45.6 g Si
0.120 kg Ti
4.79 x 1024 atoms
6.02 x 1024 atoms Bi
1.00 x 1024 atoms Mn
3.40 x 1022 atoms He
1.50 x 1015 atoms N
1.50 x 1015 atoms U
2.09 x 103 g
6.68 x 1020 atoms
91.3 g
3.45 x 1022atoms
0.226 g
9.77 x 1023atoms
3.49 x 10-8 g
1.51 x 1024 atoms
5.93 x 10-7 g

34. Practice WS #4 Two-Step Mole Conversions
WS #5 Mixed Problems w/ Naming 
WS #6 Mixed Problems

35. Percent Composition by Mass
The percent composition by mass tells you what percent an element is out of a compound.
For example:
Hydrogen (H) is 11.2% of H2O by mass

36. The Steps: Percent by Mass
Find the molar mass of the elements in the compound.
Find the molar mass of the compound.
Divided the molar mass of each element by the molar mass of the compound and multiply by 100%.
(Make sure you multiply the mass of the element by the number of moles present in the compound)

37. Example:
Find the percent composition of each of the elements in NaHCO3

38. Practice
Determine the percent by mass of each element in calcium chloride (CaCl2).
Calculate the percent by mass of each element in sodium sulfate (Na2SO4)
Calculate the percent composition of phosphoric acid (H3PO4).

39. Practice
Determine the percent by mass of each element in calcium chloride (CaCl2).
Ca = 36.11%, Cl = 63.89%
Calculate the percent by mass of each element in sodium sulfate (Na2SO4)
Na = 32.37%, S = 22.58%, O = 45.05%
Calculate the percent composition of phosphoric acid (H3PO4).
H= 3.08%, P = 31.61%, O = 65.31%

40. Practice: Percent Composition
WS #7 Percent Composition by Mass
WS #8 Percent Composition Hard

41. Molecular Formula vs. Empirical Formula
A molecular formula indicates the actual number of elements in a covalent compound.
We call the mass (in amu) of the molecular formula the molecular mass.
Remember: we can convert molecular mass to molar mass by simply changing the units (amu  g/mol).
The empirical formula for a molecule is the smallest whole-number ratio of the elements. It may or may not be the same as the molecular formula.
We call the mass (in amu) of the empirical formula the formula mass.
Remember: we can convert the formula mass to a molar mass simply by changing the units (amu  g/mol).

42. A Few Key Conceptual Points:
The molecular formula can be the same as or different from the empirical formula.
Ionic compounds ONLY have empirical formulas (lowest whole number ratios) because of their crystal lattice structure.
Only covalent compounds can have molecular formulas that differ from their empirical formulas.
For covalent compounds, it is the molecular formula that occurs in real life.

43. Example: Empirical vs. Molecular
The molecular formula for glucose (a product of photosynthesis) is C6H12O6.
The empirical formula is CH2O.
The molecular formula for glucose is 6 times the empirical formula.

44. 2 New Skills to Learn
How to calculate empirical formula from percent composition.
How to calculate molecular formula from empirical formula and molar mass.

45. #1 Calculating Empirical Formula from Percent Composition
Assume you have a 100 g sample and convert your percentages to masses.
Divide the masses that you have from step #1 by their molar masses to determine the mole ratios.
Calculate the simplest ratio by dividing each mole amount from #2 by the smallest mole amount that you calculated from #2.
Convert to the smallest whole-number ratio by multiplying all molar amounts by a common multiple. (No parts of atoms rule).

46. Example: Calculating Empirical Formula from Percent Composition
Example: Imagine that you know the percent composition of a compound, ethane.
Problem: Determine the empirical formula from the percent composition:
C = 79.9%
H = 20.1%

47. The Steps Assume you have 100 g
79.9% C = 79.9g C and % H = 20.1g H
Divide the mass of each element by its molar mass.
(79.9g ÷ g/mol) = 6.65 mol
(20.1g ÷ 1.01g/mol) = mol
Divide the molar amounts from step 2 by the smallest molar amount (6.65 mol).
(6.65 mol C ÷ 6.65 mol) = 1 C
(19.9 mol H ÷ 6.65 mol) = 3 H
Determine the ratio of elements and convert to whole numbers by multiplying by a common multiple if necessary.
Therefore, the empirical formula for ethane is CH3

48. Practice: Empirical Formula from Percent Composition
Determine the empirical formula for the following compounds based on the percent composition given.
36.84% nitrogen, 63.16% oxygen
35.98% aluminum, 64.02% sulfur
81.82% carbon, 18.18% hydrogen
N2O3
Al2S3
C3H8

49. Skill 2: Molecular Formula from Empirical Formula
A covalent compounds molecular formula is some integer multiple of the empirical formula.
i.e., Molecular formula = n(empirical formula)
Where n= natural numbers (1, 2, 3, etc.)
The molecular formula is found by comparing the ratio of the molecular mass to the formula mass.

50. Example Molecular mass of acetylene = 26.04 amu
Formula mass of acetylene = amu
The ratio of molecular mass to formula mass:
(26.04 amu)/(13.02 amu) = 2
Therefore: the molecular mass is 2x the formula mass, so the molecular formula has 2x the number of each elements found in the empirical formula.

51. The Steps
Note: You will be given the molecular mass!! Otherwise, you can’t figure this out!!
Determine the empirical formula of the compound.
Determine the formula mass of the empirical formula.
Divide the molecular mass by the formula mass.
(molar mass) ÷ (formula mass) = ????
Multiply the subscripts of each element in the empirical formula by the number determined in step 3. You’re done.

52. Example:
Determine the molecular formula for succinic acid, which is 40.68% C, 5.08% H, and 54.24% O. It has an experimentally determined molecular mass of amu.
Find the empirical formula.
C2H3O2
Calculate the formula mass of the empirical formula.
59.05 amu
Divide the molecular mass by the formula mass.
(118.1 amu) ÷ (59.05 amu) = 2
Multiply the empirical formula by the whole number you found in step 3.
2(C2H3O2) = C4H6O4

53. Practice Problems
A compound with a molar mass of g/mol is 46.68% N and 53.32% O. What is its molecular formula?
The experimentally determined molar mass of an unknown compound is g/mol, and it is 65.45% C, 5.45% H, and 29.09% O. What is its molecular formula?
N2O2
C6H6O2

54. Practice WS #9 Empirical & Molecular Formula

55. Hydrates
Some salts have the ability to bind water molecules within their lattice structure. These compounds are known as hydrated salts or hydrates, written as CoCl2 • 6H2O. The salt with no water in the crystal lattice is called an anhydrous salt CoCl2.
Heating a hydrated salt will drive off the water molecules drying out the crystal and forming the anhydrous salt.
Anhydrous salts will absorb water from the air. They are used to keep products dry by inserting little packets into shoe boxes, cameras etc.

56. Formula For a Hydrate
In the formula for a hydrate, the number of water molecules associated with each formula unit of the compound is written following a dot.
Example:
Sodium carbonate decahydrate
Na2CO3 ∙ 10H2O
The number of water molecules associated with the formula unit in the compound are indicated in the prefix.
The mass of water associated must always be included in the molar mass

57. % Composition of Hydrates
% Composition of hydrate tells you how much of the hydrate is salt, and how much is water.
It’s just like the other percent composition calculations.
Calculate the molar masses for the anhydrous salt, the water and the total mass.
Divide the molar mass of each part (salt, water) by the total molar mass and multiply by 100%. Done.
% = Part / whole X 100
% H2O = (mass H2O) ÷ (total mass) x 100
% salt = (mass salt) ÷ (total mass) x 100

58. Practice Problem What is the % composition of CaSO4 · 2H2O? H20 = 21%

59. Analyzing a Hydrate
To determine the formula for a hydrate you must determine the moles of water associated with each mole of hydrate.
The water is boiled off, and the before and after masses of the compound are compared. The difference is the mass of the water in the compound.
The differences in grams are converted to molar masses and the ratio of the formula unit to water is determined.

60. Hydrate Example
You start with 5.00 g of BaCl2 · xH2O. After heating there is only 4.26 g of BaCl2. What is the formula for the hydrate?
Find the difference between the two masses. This is the mass of water.
5.00g – 4.26g = 0.74g H2O
Divide the mass of each (salt, water) by their molar masses. This gives you the # of moles.
BaCl2 : (4.26g)/(208.23g/mol) = mol BaCl2
H2O: (0.74g)/(18.02g/mol) = mol H2O
Divide the moles of water by the moles of salt. This gives you the mole ratio of the hydrate. You’re done.
(0.041 mol)/( mol) = 2/1 (Note: if you get decimals, round).
So there are 2 moles of water per mole of barium chloride
Barium chloride dihydrate = BaCl2 ∙ 2H2O

61. Practice Problems
During a lab, 1.62g of CoCl2 · xH2O is heated. After heating, 0.88g CoCl2 remains. What is the formula for the original hydrate?
CoCl2 · 6H2O

62. Practice WS #11 Hydrate Problems Mole PRACTICE Quiz Mole Quiz
Mole Test Review