1. Warm up Give the units for the following measurements
Length----- Meters
Mass---- Grams
Volume--- Liters
Temperature--- Celsius
Density ---g/ml

2. Unit 6 – Chemical Quantities & The Mole
Mole Math Notes
Day 1

3. Day 1: Mole Math
We use a unit called the , or , to measure the amount of a substance. The mole can represent or
mole
mol
mass (grams),
volume (liters),
amount (particles).

4. **You must ALWAYS go through the MOLE!**
Avogadro's number: and is also called a .
One dozen = _________ items.
3 dozen cookies = _______ cookies
0.5 dozen doughnuts = _________ doughnuts
A “dozen” is a counting unit equal to of any object.
A “Mole” is a counting unit equal to of any object, even really small ones like
molar mass or atomic mass
22.4 Liters
6.02 x 1023 particles
6.02 x 1023 particles
Mole 12 36 6 12
6.02 x 1023
atoms,
molecules,
or formula units.

5. What is so special about 6. 02 x 1023
What is so special about 6.02 x 1023 ? Why do scientists use that number?
Conversion factor
Allows us to manipulate many small particles, as if they were one whole part.
Example Conversion Factors:
1 mole = 6.02 x 1023 particles (atoms, molecules, or formula units)
1 mole of copper = 6.02 x of copper.
1 mole of CuCl2 = formula units of CuCl2
1 mole of CO2 = molecules of CO2
The amount of a substance containing Avogadro’s number of any kind of chemical unit is called a of that substance.
One mole of K contains atoms.
One mole of NaOH contains formula units.
atoms
6.02 x 1023
6.02 x 1023
mole
6.02 x 1023
6.02 x 1023

6. Molar Mass #1 One atomic mass 28.09 g Si 1 1 mol Si = __________
For any Monatomic element _________mole of
that element is equal to its ________ __ Example:
1 mol Si = __________
40.08g Calcium = ___________ mol Ca
(its atomic mass from the periodic table = 1 mole)
Get out your periodic table : )
atomic mass
28.09 g Si 1

8. Diatomic Elements never exist alone in nature!!!
*Trick for remembering the 7 diatomics *

9. Molar Mass
Molar mass is the mass of
of a substance.
Other names for molar mass include…
one mole
*formula mass
*gram formula mass

10. Molar Mass #1 mole atomic mass 12.011 28.014
One _____________of any element will have a mass in grams corresponding to the value of its ____________________.
1 mol carbon = __________ g/mol
1 mol nitrogen = ___________ g/mol
(One mole of oxygen has a mass of g/mol-
We will Use g/moles for Oxygen
mole
atomic mass
12.011
28.014

11. Molar Mass of Compounds
One ________ of any molecule/compound will have a mass in grams corresponding to the value of its molar mass (the sum of the masses of the elements that compose it).

12. Example: Water, H2O H O 1.01 16.00 X 2 1 2.02 g 16.00 g 18.02 g/mol
Element Molar Mass # of Atoms Total H O
1.01
16.00 X 2 1
2.02 g
16.00 g
18.02 g/mol
This is the mass of 1 mol of water!

13. Ex. Ca(NO3)2
Ca: x 1 =
N: x 2 =
O: x 6 =
g/mole

14. Sample Problems H2CO3 62.03 g/mol H2CO3 1.01 2 2.02 12.01 1 12.01
Carbonic Acid ___________
# of H atoms: x = +
# of C atoms: x =
# of O atoms: x =
___________________
1.01 2
2.02
12.01 1
12.01
16.00 3
48.00
62.03 g/mol H2CO3

15. Sample Problems (NH4)2SO4 132.17 g/mol (NH4)2SO4 14.01 2 28.02 1.01 8
Ammonium sulfate  _____________
# of N atoms: x = +
# of H atoms: x =
# of S atoms: x =
# of O atoms: x =
_________________________
14.01 2
28.02
1.01 8
8.08
32.07 1
32.07
16.00 4
64.00
g/mol (NH4)2SO4

16. Practice HW 1. What is the atomic mass of sodium?
2. Calculate the molar mass of Al2(SO4)3
3. Calculate the molar mass of nitrogen (hint: diatomic!?!?).

17. Practice 1. What is the atomic mass of sodium?
23.00 g/molNa or g/molNa
2. Calculate the molar mass of Al2(SO4)3
342.17 g/mol Al2(SO4)3
Al 2 x = 53.96
S 3 x =96.21
O 12x 16.00=
3. Calculate the molar mass of nitrogen (hint: diatomic!?!?).N 2 2x 14.01= g/mol

18. Practice 4. H2SO4? g/mol 5. Calculate the molar mass of Ethanol C2H5OH
Calculate the molar mass of Potassium Chloride g/mol
Copper II Sulfate g/mol

19. Practice
4. H2SO4? g/mol H2SO4
5. Calculate the molar mass of Ethanol C2H5OH
46.07 g/mol C2H5OH
Calculate the molar mass of
Potassium Chloride KCl g/mol KCl
7. Copper II Sulfate CuSO g/mol CuSO4
Cu—63.55 S
O---4 x 16= 64
159.62

20. Homework Molar Mass
Show all work and units to receive full credit.

21. Mole Highway
NOTES:
To convert between units, follow the highway. Notice, there is no shortcut from grams to liters or between any of the three units surrounding the mole.
This means you have to convert to moles before converting to another unit!

22. Draw the Mole Road Map!!

23. If you have grams and want moles
Gram of element x 1 mole of element = Mole of element
Molar mass of element

24. Mole Humor
What was Avogadro’s favorite board game?
Moleopoly

25. Molar mass:
called gram atomic mass when single element is used.
called gram formula unit when ionic compound is used.
called gram molecular unit when molecular compound or diatomic molecules used.
Diatomic molecules are atoms that bond with themselves. There are SEVEN of these that you need to remember: Br2 I2 N2 Cl2 H2 O2 F2 (Remember this by the name “BRINClHOF”)

26. Determine the number of moles in….
25 g sodium
1 mol Na
22.99g Na
25 g Na
= 1.1 mol Na
Molar mass of Sodium =
Na = 1 x = g

27. Determine the number of moles in….
85 g H2SO4
85 g H2SO4
1 mol H2SO4
g H2SO4
= mol H2SO4
Mass in grams of H2SO4 =
H = 2 x = 2.016
S = 1 x =
O = 4 x =
Add to get molar mass of g

28. Determine the number of grams in….
2.5 moles of sodium
2.5 mol Na
22.99 g Na
I mol Na
= 57 g Na
Molar mass of Sodium =
Na = 1 x = g

29. Determine the number of grams in….
0.50 moles of H2SO4
0.50 mol H2SO4
g H2SO4
I mol H2SO4
= 49 g H2SO4
Mass in grams of H2SO4 =
H = 2 x =
S = 1 x =
O = 4 x =
Add to get molar mass of g

30. How many moles are there in 27 g of ethanol (C2H5OH)?
Mass of C2H5OH:
C = 2 x =
H = 6 x 1.008=
O = 1 x =
Add these all up for mass in grams = g

31. Homework
Show all work and units to receive full credit.

32. Representative Particles  Moles

33. It’s a quantity, and it’s a BIG one.
6.02 x 1023
Avogadro’s Number = ______________ and is also called a _______.
Like a “pair” or “dozen,” a “mole” represents a set number of things; in chemistry, those “things” are particles.
For example:
1 dozen = ______ items
3 dozen = ______ cookies
.5 dozen = ______ doughnuts
Mole 12 36 6

34. How big is a mole?

35. HOW LARGE IS IT???
1 mole of hockey pucks would equal the mass of the moon!
1 mole of basketballs would fill a bag the size of the earth!
1 mole of pennies would cover the Earth 1/4 mile deep!

36. It’s a quantity, and it’s a BIG one.
So, a “dozen” is a counting unit equal to __12__of any object. Likewise, a “Mole” is a counting unit equal to _____________ of an object, even really small ones like ____________, ____________, or _____________.
6.02 x 1023
atoms
molecules
formula units

37. What’s so special about 6.02 x 1023?
Why do scientists use that number?
A conversion factor.
Allows us to manipulate many small particles, as if they were one whole part.
For elements on the periodic table, there is a 1 to 1 relationship between the mass of a single atom (in amu) and the mass of 1 Mole of the same species of atom (in grams).
Remember, Mole is a quantity…6.02 x 1023 particles.

38. What’s so special about 6.02 x 1023?
Atomic Number
(number of protons)
Using the periodic table, we’ve learned that an Argon atom has a mass of amu. Unfortunately, manipulating a single atom of any element isn’t reasonable, so taking its individual mass isn’t possible. However, we do have the Mole…
This is the MOLAR MASS, or mass of a mole of a substance.
Atomic Mass
(average number of protons and neutrons)
So, a g sample of Argon contains 6.02 x 1023 atoms…that’s Avogadro’s Number, the MOLE!
It’s a 1 to 1 relationship between a single atom in amu and a Mole of atoms in grams. Ar
39.95 amu
39.95 g
1 Atom of Argon
1 Mole of Argon

39. Representative Particles
A representative particle is the smallest unit of a substance.

40. Types of Particles atoms molecules formula units molecules atoms
Monatomic elements = _______________
Diatomic elements = _________________
Ionic compounds = _________________
Covalent compounds = ______________
Ions = _________________
Acids = _________________
atoms
molecules
formula units
molecules
atoms
molecules

41. Avogadro’s number, which is ________________, represents the number of “chemical units” in one mole of any substance. For the monatomic elements, the “chemical unit” is an __________. 1 mol of any chemical = ______________ particles.
6.02 x 1023
atom
6.02 x 1023

42. Examples 6.02 x 1023 formula units 6.02 x 1023 atoms
1 mol CaCl2 =
1 mol Ca2+ =
1 mol HCl (aq) =
1 mol P2O5 =
1 mol Ca =
1 mol Cl2 =
6.02 x 1023 formula units
6.02 x 1023 atoms
6.02 x 1023 molecules
6.02 x 1023 molecules
6.02 x 1023 atoms
6.02 x 1023 molecules

44. Sample problems = 74.8 moles Mn 1 mol Mn
How many moles are in 4.50 x 105 atoms of manganese?
4.50 1025 atoms Mn
1 mol Mn
6.02  1023 atoms Mn
= 74.8 moles Mn

46. Sample problems 3.27 mol 6.02  1023 atoms 1 mol = 1.97  1024 atoms
2. How many atoms are found in 3.27 mol of magnesium?
3.27 mol
6.02  1023 atoms
1 mol
= 1.97  1024 atoms

48. Sample problems 3.4 mol 6.02  1023 formula units 1 mol
3. Chalk is composed primarily of calcium carbonate. How many particles are in 3.4 moles of calcium carbonate?
3.4 mol
6.02  1023 formula units
1 mol
= 2.0  1024 formula units

49. Moles of Chalk Lab

50. Homework
Show all work and units to receive full credit.

51. Notes: Volume <-> Moles
Warm-up:
1. How many atoms are in 0.75 mol of zinc?

55. 2. At STP, how many moles are found in 54L of neon gas?
Sample Problems
(SHOW ALL WORK & UNITS to receive full credit.)
1. The average lung capacity of a male is 6.0L. The average lung capacity of a female is 4.7L. Assume the following:
If your TEACHER’s lungs are completely filled with oxygen, determine the number of moles of oxygen gas in the lungs of your chemistry teacher at STP.
2. At STP, how many moles are found in 54L of neon gas?
3. How many liters are found in 3.02mol of helium at STP?

56. Homework page 12

57. Empirical vs. molecular

58. Empirical formula Is the simplest form of a formula
Is written in the lowest possible ratio
E.g., CH2O or CH
Will be reduced. May or may not exist in this form in the real world.

59. Molecular formula
Is the true formula or actual ratio of the atoms in a compound.
E.g., C2H4O2 or C6H6
Will not be reduced. Formula describes a substance as it actually exists.

60. Empirical or Molecular?
Na2O ___________________
C3H6 ___________________
K2SO4 ___________________
C6H12O6 ___________________
empirical & molecular
molecular
empirical & molecular
molecular
Keep in mind that for some compounds, its empirical formula can also be its molecular formula.

61. Which pair has the same empirical formula?
Na2O and Na2O2
C6H12O6 and CH2O
C3H6 and C5H12
C6H6 and C5H5

62. Calculating an Empirical Formula
Determine the mass of each element.
Convert the mass of each to moles. 
Find the mole to mole ratios of each element by dividing the number of moles of each element by the smallest number of moles.
If the ratio is not a whole number, multiply each ratio by a factor to make them all whole numbers.
Write the formula using the mole ratio as the subscript for the formula.

63. Find the empirical formula of 69.5% O & 30.5% N.
Step 1: Divide % or grams by its atomic mass to get moles of each element.
69.5 g O
1 mol O
g O
= mole O
25.9 g N
1 mol N
g N
= mole N

64. Find the empirical formula of 69.5% O & 30.5% N.
Step 2: Divide smallest mole number in each element to get ratio of that element.
4.344 mole O
= 2 O
2.177 mole
2.177 mole N
= 1 N
2.177 mole

65. Find the empirical formula of 69.5% O & 30.5% N.
*This answer becomes the subscript for that element; round to nearest whole number if .8 or higher or lower than .2.
Answer = NO2

66. Independent Practice:
1. Analysis of a compound shows that it contains 10.88g of calcium and 19.07g of chlorine. Determine the empirical formula of this compound.

67. Independent Practice:
2. One of the most commonly used white pigments in paint is a compound of titanium and oxygen that contains 5.99 g titanium by mass and 4.01 g oxygen by mass. Determine the empirical formula and name for this compound.

68. Independent Practice:
3. Used in the production of nylon, adipic acid is an organic compound composed of 49.31% C, 43.79% O, and the rest is hydrogen. Determine the empirical formula of adipic acid.

69. Calculating Molecular Formulas
The molecular formula will have the same ratio as the empirical formula.
To determine a molecular formula, we will multiply the empirical formula by a whole number factor (WNF).

70. Calculating Molecular Formulas

71. (P2O5)2  P4O10 empirical mass = 141.943 g/mol 283.88 g/mol
1. Empirical Formula = P2O5
Molar Mass= g/mol
What is the molecular formula of this compound?
empirical mass = g/mol
g/mol
g/mol
WNF =
= 2
(P2O5)2  P4O10

72. 2. Nitrogen and oxygen form multiple molecular compounds together
2. Nitrogen and oxygen form multiple molecular compounds together. One of these compounds is used to fuel space shuttles and has the empirical formula NO2. If the molar mass of this compound is g/ mol, what is the molecular formula?

73. **Fix typo in your packet, please.
3.Butane is commonly used in lighters. It is composed of 17.37% hydrogen and 82.63% carbon . It has a molar mass of g/mol. What is the molecular formula of butane?
**Fix typo in your packet, please.

74. 4. Vitamin C is 40. 91% C, 4. 587% H, and the remaining is oxygen
4. Vitamin C is 40.91% C, 4.587% H, and the remaining is oxygen. If the molar mass of Vitamin C is about 180 g/mol, determine the empirical and molecular formula.

75. Complete pages 20-21 for Homework!!!!