1. Section 3.5—Counting Molecules
So the number of molecules affects pressure of an airbag…how do we “count” molecules?

2. What is a Mole? Ted Ed video

3. What is a mole? Mole – metric unit for counting
We use it just like we use the terms dozen and ream!
The only acceptable abbreviation for “mole” is “mol”…not “m”!!

4. What is a counting unit?
You’re already familiar with one counting unit…a “dozen”
A dozen = 12
“Dozen” 12
A dozen doughnuts
12 doughnuts
A dozen books
12 books
A dozen cars
12 cars

5. What can’t we count atoms in “dozens”?
Atoms and molecules are extremely small
We use the MOLE to count particles

6. A mole = 6.02  1023 particles (called Avogadro’s number)
6.02  1023 = 602,000,000,000,000,000,000,000
“mole”
6.02  1023
1 mole of doughnuts
6.02  1023 doughnuts
1 mole of atoms
6.02  1023 atoms
1 mole of molecules
6.02  1023 molecules
This number was named after Amadeo Avogadro. He did not calculate it!

7. FUNNY!

8. Representative Particles
Remember, matter is broken down into either SUBSTANCES or mixtures
Substances are broken down into either ELEMENTS or COMPOUNDS
Type of Matter
Example
Representative Particle
Element Fe
Atom
Ionic Compound
NaCl
Formula Unit
Covalent Compound
CO2
Molecule

9. How many molecules of water are in 1.25 moles?
Example: Particles & Moles Use the conversion factor (1 mol = 6.02 x 1023) particles to convert
Example 1:
How many molecules of water are in 1.25 moles?

10. Example: Molecules & Moles
How many molecules of water are in 1.25 moles?
1 mol = 6.021023 molecules
1.25 mol
H2O
Molecules
H2O
6.02  1023
= _______ molecules H2O
7.531023 1
mol H2O

11. How many moles are equal to 2.8 × 1022 formula units of KBr?
Let’s Practice #2
Example:
How many moles are equal to 2.8 × 1022 formula units of KBr?

12. How many moles are equal to 2.8 × 1022 formula units KBr?
Let’s Practice #2
Example:
How many moles are equal to 2.8 × 1022 formula units KBr?
1 mol = 6.021023 formula units
2.8 × 1022 formula units 1
mole
= _______ moles
0.047
6.02  1023
Formula units

13. How many atoms are equal to 3.56 moles of Fe?
Let’s Practice #3
Example:
How many atoms are equal to moles of Fe?

14. How many atoms are equal to 3.56 moles of Fe?
Let’s Practice #3
Example:
How many atoms are equal to moles of Fe?
1 mol = 6.021023 molecules
3.56 moles Fe
6.02 x 10 23
atoms
= _______ atoms
2.14 x 1024 1
moles

15. Molar Mass Molar Mass – The mass for one mole of an atom or molecule.
Other terms commonly used for the same meaning:
Molecular Weight
Molecular Mass
Formula Weight
Formula Mass

16. Molar Mass for Elements
The average atomic mass = grams for 1 mole
Average atomic mass is found on the periodic table
Element
Mass
1 mole of carbon atoms (C)
12.01 g
1 mole of oxygen atoms (O2)
16.00 g x 2 = g O2
1 mole of hydrogen atoms (H2)
1.01 g x 2 = 2.02 g H2
Unit for molar mass: g/mole or g/mol

17. Molar Mass for Compounds
The molar mass for a molecule = the sum of the molar masses of all the atoms

18. Calculating a Molecule’s Mass
To find the molar mass of a molecule: 1
Count the number of each type of atom 2
Find the molar mass of each atom on the periodic table 3
Multiply the # of atoms by the molar mass for each atom 4
Find the sum of all the masses

19. Find the molar mass for CaBr2
Example: Molar Mass
Example:
Find the molar mass for CaBr2

20. Find the molar mass for CaBr2
Example: Molar Mass 1
Count the number of each type of atom
Example:
Find the molar mass for CaBr2 Ca 1 Br 2

21. Find the molar mass for CaBr2
Example: Molar Mass 2
Find the molar mass of each atom on the periodic table
Example:
Find the molar mass for CaBr2 Ca 1
40.08 g/mole Br 2
79.90 g/mole

22. Find the molar mass for CaBr2
Example: Molar Mass 3
Multiple the # of atoms  molar mass for each atom
Example:
Find the molar mass for CaBr2 Ca 1 
40.08 g/mole =
40.08 g/mole Br 2 
79.90 g/mole =
g/mole

23. Find the molar mass for CaBr2
Example: Molar Mass 4
Find the sum of all the masses
Example:
Find the molar mass for CaBr2 Ca 1 
40.08 g/mole =
40.08 g/mole Br 2 
79.91 g/mole = +
g/mole
g/mole
1 mole of CaBr2 = g

24. Example 2: If you see a Parentheses in the Formula
Be sure to distribute the subscript outside the parenthesis to each element inside the parenthesis.
Example:
Find the molar mass for Sr(NO3)2

25. Example 2: Molar Mass & Parenthesis
Be sure to distribute the subscript outside the parenthesis to each element inside the parenthesis.
Example:
Find the molar mass for Sr(NO3)2 Sr 1 
87.62 g/mole =
87.62 g/mole N 2 
14.01 g/mole =
28.02 g/mole O 6 
16.00 g/mole = +
96.00 g/mole
g/mole
1 mole of Sr(NO3)2 = g

26. Find the molar mass for Al(OH)3
Let’s Practice #3
Example:
Find the molar mass for Al(OH)3

27. Find the molar mass for Al(OH)3
Let’s Practice #2
Be sure to distribute the subscript outside the parenthesis to each element inside the parenthesis.
Example:
Find the molar mass for Al(OH)3 Al 1 
26.98 g/mole =
26.98 g/mole O 3 
16.00 g/mole =
48.00 g/mole H 3 
1.01 g/mole = +
3.03 g/mole
78.01 g/mole
1 mole of Al(OH)3 =78.01 g

28. Using Molar Mass in Conversions
What is molar mass?

29. Example: Moles to Grams
How many grams are in 1.25 moles of water?

30. Example: Moles to Grams
When converting between grams and moles, the molar mass is needed
Example:
How many grams are in 1.25 moles of water? H O 2 1
1.01 g/mole
16.00 g/mole  =
2.02 g/mole +
18.02 g/mole
1 mole H2O molecules = g
1.25 mol H2O
18.02
g H2O
= _______ g H2O
22.5 1
mol H2O

31. Example: Grams to Moles
How many moles are in
25.5 g NaCl?

32. Example: Grams to Moles
How many moles are in
25.5 g NaCl Na Cl 1
22.99 g/mole
35.45 g/mole  = +
58.44 g/mole
1 moles NaCl molecules = g
25.5 g NaCl 1
mol NaCl
58.44
g NaCl
= ____ moles NaCl
.436

33. Example: Grams to Molecules
How many formula units are in
25.5 g NaCl?

34. Example: Grams to Moles
How many formula units are in
25.5 g NaCl Na Cl 1
22.99 g/mole
35.45 g/mole  = +
58.44 g/mole
1 moles NaCl formula units = g
25.5 g NaCl 1
mol NaCl
6.02 x 1023 FU’s
58.44
g NaCl
1 mol NaCl
2.63 x 1023
= ____ FU’s NaCl

35. Gases & Moles
Amounts

36. Molar Volume The molar volume of a gas is measured at
STP (standard temperature and pressure)
1 mole of any gas = 22.4 L

37. Molar Volume as a Conversion Factor
The molar volume at STP
has about the same volume as 3 basketballs
can be used to form 2 conversion factors:
22.4 L and mole
1 mole L

38. Let’s Try it Out! Example
When converting between volume and moles, STP must be a condition to use molar volume
Example:
An experiment
requires .0580
moles of NO.
What volume
would you need
at STP?
1 mole NO = 22.4 L
.O580 mol NO
22.4
L NO
= _______ L NO
1.2992 1
mol NO

39. Suppose you need 4.22 g of Cl2. What volume at STP would you use?
Try Another!
Example:
Suppose you need 4.22 g of Cl2. What volume at STP would you use?
1 moles Cl2 = g Cl2
1 mol = 22.4 L at STP
4.22 g Cl2 1
L Cl2
mol Cl2
22.4
70.90 g Cl2 1
mol Cl2
= _________ L Cl2
1.33

40. Percent Composition
Defined as the percent by mass of each element in a compound
Steps to Finding Percent Composition
Add up the mass of each element within the compound to get the mass of the compound.
Divide each element’s mass by the mass of the compound.
Multiply by 100
% composition= mass of element x 100
mass of compound

41. Unit 3.6 Formula Calculations

42. Percent Composition by Mass of Air

43. Example: Calculate the % composition of each element in calcium carbonate.
CaCO3
Molar mass = g
% C = 12.01/ x 100 = %
%Ca = 40.08/ x 100 = 40.04%
%O = 48.00/ x 100 = 47.96%

44. Example: What is the % of each element in a compound that contains 29
Example: What is the % of each element in a compound that contains 29.00g Ag and 4.30g S only?
Total mass of compound = g
% Ag = 29.00/33.30 x 100 = %
%S = 4.30/33.30 x 100 = 12.9%

45. Hydrates A HYDRATE is an ionic compound with water trapped
in its crystal.
Examples are:
CuSO4 5H2O MgSO4 7 H2O CoCl2 H2O
Heating a hydrate removes the water and leaves
behind just the salt which is called the anhydrate.

46. Example: What is the % water in the hydrate, CuCl2  2H2O
Molar mass of hydrate = g
% water = 36.04/ x 100 = 21.14%

47. Heating of A Hydrate Animation
Calculating the experimental % composition of water in a hydrate.

48. Empirical Formula
A chemical formula showing the simplest whole number ratio of moles of elements (subscipts)
Hydrogen Peroxide has an actual formula (molecular formula) of H2O but an
empirical formula of HO

49. How to Calculate Empirical Formula
RHYME: Percent to Mass
Mass to Mole
Divide by Small
Multiply til Whole
Assume 100 grams of the sample of compound. Switch the percent sign to grams
Convert each element’s mass into moles.
Divide each element’s mole amount by the smallest mole amount in the entire problem. The answer is the subscript of the element within the compound.
OPTIONAL: If mole ratio is not within .1 of a whole number, multiply each amount by the smallest whole number that will produce either a whole number itself or a number within .1 of a whole number.

50. Example: What is the empirical formula for 40.05% S and 59.95% O?
Switch the percent sign to grams & convert each element’s mass into moles
40.05 g S / 32.01g = mol S
59.95 g O / g = mol O
Divide each element’s mole amount by the smallest mole amount in the entire problem.
1.250 mol S = mol O = 2.99 = 3
1.250 mole mol
S1O3  SO3

51. Example: What is the empirical formula for 43.64% P and 56.36% O?
Switch the percent sign to grams & convert each element’s mass into moles
43.64 g P / 30.97g = mol S
56.36 g O / g = mol O
Divide each element’s mole amount by the smallest mole amount in the entire problem.
1.409 mol S = mol O = 2.49 ≠ 3
1.409 mole mol
If mole ratio is not within .1 of a whole number, multiply each amount by the smallest whole number that will produce either a whole number itself or a number within .1 of a whole number.
1 x 2 = x 2 = = 5
P2O5

52. Molecular Formula Is the ACTUAL, true formula of the compound.
They are usually multiples of their empirical formula
N2O4 is the molecular formula; the empirical formula is NO2
Notice that the molecular formula is 2 times larger than the empirical formula

53. Molecular Formula

54. How to Calculate Molecular Formula
1. You need to find the empirical formula and calculate its molar mass. Call this empirical formula mass EFM.
2. Find the mass of the actual formula which will most likely be given to you in grams. Call this molecular formula mass MFM.
3. Divide the MFM by the EFM to get a factor.
4. Multiply the factor by the empirical formula to get the MOLECULAR FORMULA

55. Example: What is the molecular formula of a compound whose empirical formula is CH4N and the molecular mass is g/mol?
1. Empirical Formula Mass (EFM) = = g
2. Molecular Formula Mass (MFM) = g
/ = 2
4. 2(CH4N) = C2H8N2