1. Measurements of Matter
Mass, volume, length, density… all these are measurements of matter.
What is matter made of though? atoms
If you measured the density of an object to be 2.04 g/mL, you have a certain number of atoms all joined together in a certain amount of space.
So how many atoms are in there?

2. The MOLE It’s a beauty mark… It’s a small furry garden pest…
No, wait… its how we count ATOMS!

3. What is a Mole and how is it used in Chemistry?
A mole is a quantity of atoms. Since atoms are so tiny they are hard to individually measure, so scientist measure them in large quantities.
Problem: How do you count something as small as an atom?

4. Solution: the mole
Specific quantity (amount) of atoms in a substance. Basically, a word that represents a specific number (ex: pair = 2)
1 mole = 6.02 x representative particles (a lot of very tiny particles)
This value is called Avogadro’s Number

5. What are representative particles?
smallest particle that retains chemical and physical properties
3 types depending on the compound:
Atoms: Single element
Molecules: covalent compound
Formula Units: Ionic Compounds or ions

6. What is a mole?
Try this

7. How do we use it?
Used to count groups of small particles for the purpose of mixing them with other particles in the lab.
Ex: If 2.5 moles of sodium were used in an experiment, how many atoms of sodium were there?
1 mole = 6.02x1023
2.5 mole = ___________

8. Ex: Exactly x 1019 molecules of water are needed to safely react in with sodium. How many moles of water are needed?

9. The Mole
Used to count groups of small particles for the purpose of mixing them with other particles in the lab.
Problem – no way to count a mole of anything. (atoms are too small to see) If we can’t count them – must be another way to measure them
Solution – measure the mass of a mole with a balance (molar mass)

10. Using Molar Mass as a conversion factor
The periodic table contains the molar mass of all elements.
Molar Mass = mass of 1 mole (6.02x1023)of its atoms.
6.02x1023 atoms Carbon = 1 mole carbon = grams Carbon
If 2.5 moles of carbon were used in an experiment, how many grams of carbon must you weigh out?

11. Using molar mass
How many grams are in 15.7 mole MgCl2

12. The Mole of a gas
Why would it be difficult to measure the mass of helium?
(how would it stay on the balance)
Scientist have discovered that 1 mole of any gas has a constant volume at STP. (standard temperature & pressure)
1 mole = 22.4 L of any gas
This value is called Molar Volume

13. Using molar volume as a conversion factor
How many moles are in L of oxygen gas?
How many liters are in moles of nitrogen gas?

14. Number of particles – Avagadro’s number
What does a mole tell us?
Number of particles – Avagadro’s number
Mass particles in solid or liquid – Molar mass
Volume of particles in a gas – Molar Volume

15. Molar Mass
The mass of a substance.

16. The subscript is the number at the bottom of a formula.
FORMULAS - review
MgCl2
The subscript is the number at the bottom of a formula.
There is 1- Mg & 2 - Cl

17. Calculate the molar mass
Calculate the molar mass of 1 mole of magnesium chloride. (first you need a formula)

18. How to calculate molar mass
Identify the # of atoms of each element
Multiply # atoms by the atomic mass of that element. (round to 2 #’s after the decimal)
Add them all together
Grams (g) is the unit

19. MgCl2 Mg – 1 (24.31) = 24.31 Cl – 2 (35.45) = 70.90 or 95.21 g/mol
Molar Mass
MgCl2 Mg – 1 (24.31) = Cl – 2 (35.45) = or g/mol
1 mole MgCl2 =
95.21 g MgCl2

20. 1 mole of Ammonium phosphide
Calculate the molar mass: (you MUST write the formula correctly before answering)
1 mole of Ammonium phosphide
1 mole of Trinitrogen pentachloride

21. What is a percentage mean?
How is a percent calculated?
What is the percentage of girls in the class?
What is the percentage of students in Mrs. Muchnick’s class if she has 10 girls and 12 boys?

22. Percent Composition
The percent BY MASS of each element in a compound – divide the element’s total mass (part) by the molar mass (whole) then multiple by 100 to get the percent.
Ex: % composition of MgCl2
Mg – 1 (24.31) = / x 100 = % Mg
Cl – 2 (35.45) = / x 100 = % Cl
Molar mass = g/mol
(PART)
(WHOLE)
(PART)
(WHOLE)

23. Practice - Calculate the % comp of KMnO4:
K – 1 (39.10) = / x 100 = 24.74% K Mn – 1 (54.94) = / x 100 = 34.76% Mn O – 4 (16.00) = / x 100 = 40.50% O molar mass KMnO4 =

25. Calculating the amount of an element in a sample
Find the % comp of the element in the compound
Change the % to a decimal (move decimal 2 times to the left or divide by 100)
Multiply that decimal by the amount (g) of the sample.
Ex: Calculate amount of chlorine in grams of MgCl2. (use the % we found earlier)
74.47% Cl = x = grams Cl

26. Practice: Calculate the amount of oxygen in 15.75 grams of water.
Molar mass H2O =
/ x100 = % O
88.79% O = x = grams O
Complete % comp worksheet

27. Using Molar Mass Remember – Molar mass is the mass (grams) of 1 mole
1 mole Fe = _________grams Fe
2.5 mole Fe = ________ grams Fe
113.5 grams Fe = _______ moles Fe
Mass to mole = divide by molar mass
Mole to mass = moletiply by mole mass 

28. Using molar mass
In the lab, Mrs. Mathieson needs 2.57 moles of NaCl to do an experiment. How many grams would be needed to equal moles of NaCl?
After doing the experiment, Mrs. Mathieson has 1.02 moles of NaCl remaining – how many grams does that equal?

29. Empirical Formula
The lowest whole number ratio (subscripts) of elements in a compound.
Cannot be reduced!!!
not empirical empirical
Ex: C6H12O6  CH2O

30. Molecular Formula Actual number of atoms in a chemical compound
EX: C12H24O12
Molecular Formulas can be reduced to Empirical Formulas
molecular empirical
EX: C12H24O12  CH2O
Different molecular formulas can have similar empirical formulas
molecular empirical
EX: N3O 
N12O36 
Molecular formula: C76H52O46
Empirical formula: ___________
NO3

31. C2H4 CH2 - empirical NO3 S9Cl12 C3Cl9 N4S9 S3Cl4 - empirical
PRACTICE: 1. Identify each as empirical (can’t be reduced) or molecular (can be reduced) 2. If its molecular – write the empirical
C2H4
NO3
S9Cl12
C3Cl9
N4S9
CH2 - empirical
molecular
empirical
S3Cl4 - empirical
molecular
CCl3 - empirical
molecular
empirical

32. Finding Empirical Formula from Percent Composition
Ex: A compound was found to be 54.53% Carbon, 9.15% Hydrogen, and 36.32% Oxygen. Find its Empirical Formula.
Steps:
Assume a 100g sample (change %  g)
÷ by the molar mass of that element to find moles (sig fig it!)
Identify the lowest # of moles and ÷ them all by that number
Round each to the nearest whole #
The resulting whole #are the subscripts for that element in the empirical formula

33. Calculating Empirical Formula
63.5% Silver % Nitrogen % Oxygen
63.5 g Ag 8.2 g N g O
.589 mole Ag .59 mole N mole O
AgNO3

34. Calculating Empirical Formula (special)
60.00%C 4.48%H %O
60.00g C 4.48g H g O
4.996 mole C 4.44 mole H mole N
x4 x4 x4
C9H8O4

35. Calculating Molecular Formula
Find the empirical formula
Calculate the molar mass of your empirical formula
Identify the molar mass of your molecular (GIVEN in the problem every time!)
Divide the molecular mass / empirical mass
Round to the nearest whole #
Multiply the whole # by the subscripts in the Empirical formula

36. Practice
If a compound has an empirical formula of NO3 and a molecular mass of 186g – what is the molecular formula?
Empirical formula: NO molar mass: 62.01g
Molecular mass (given) g
empirical mass
3 x NO = N3O9

37. How many f.u. are in 2.45 moles of NaCl?
How many moles are in 5.90 x 1026 molecules of CO?