1. Moles
6.022 X 1023
What is a mole?!?

2. Molar Mass The Mass of 1 mole of a substance (in grams)
Equal to the average atomic mass from periodic table (round to the tenths place)
1 mole of C atoms = g
1 mole of Mg atoms = 24.3 g
1 mole of Cu atoms = _____

3. Learning Check! =____________ 1 mole of Sn atoms 1 mole of Br2 atoms
Find the molar mass (round to the tenths place)
=____________
1 mole of Sn atoms
1 mole of Br2 atoms
=____________

4. Molar Mass of Molecules and Compounds
Formula Mass - Mass in grams of 1 mole equal to the sum of the atomic masses
1 mole of CaCl2
1 mole Ca x 40.1 g/mol = g/mol
+ 2 moles Cl x 35.5 g/mol = g/mol
111.1 g/mol CaCl2

5. How do we measure items?
You can measure mass,
or volume,
or you can count pieces.
We measure mass in grams.
We measure volume in liters.
We count pieces in MOLES.

6. We’re not talking about this kind of mole!
What is the mole?
We’re not talking about this kind of mole!

7. Moles (is abbreviated: mol)
A mole is defined as the number of carbon atoms in exactly 12 grams of carbon-12.
1 mole = 6.02 x particles.
Treat it like a very large dozen
6.02x is called: Avogadro’s number.

8. Just How Big is a Mole?
Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles.
If you had Avogadro's number of unpopped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles.
If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.
Cue song!

9. The Mole 1 dozen cookies = 12 cookies 1 mole of cookies = 6X1023
1 dozen cars =_12 cars
1 mole of cars= 6X1023 Cars
1 dozen Al atoms =________ atoms
1 mole of Al atoms =_______ atoms
Note that the NUMBER is always the same, but the MASS is very different!
Mole is abbreviated mol (gee, that’s a lot quicker to write, huh?)

10. A Mole of Particles Contains 6.02 x 1023 particles
= x 1023 C atoms
= x 1023 H2O molecules
= x 1023 NaCl “molecules”
1 mole C
1 mole H2O
1 mole NaCl

11. Avogadro’s Number as Conversion Factor
6.02 x 1023 particles
1 mol or
1 mol
Note that a particle could be an atom OR a molecule!

12. So if one mole of C = 12 g
How many moles is 24g? ____
What did you do?

13. Divided by Molar mass
So to set up always Start with what you are given then use dimensional analysis to get you to what you want.
24g C x 1mole C = 2 moles
12g
Cancel units!

14. Learning Check Show in Dimensional analysis
1. Number of grams in mole of Al
2. Number of grams in 2.0 moles of CO2

15. Chalk Lab
Weigh a piece of ordinary chalk and write your name on the sidewalk. Weigh the chalk again, and determine the number of moles of calcium carbonate that were used. Show your work using dimensional analysis. Weight of chalk before writing your name: _______________ Weight of chalk after writing your name: _______________ Grams of chalk required to write your name: _______________

16. Empirical and Molecular Formulas

17. Chemical Formulas of Compounds
Formulas give the relative numbers of atoms or moles of each element in a formula unit - always a whole number ratio (the law of definite proportions).
NO atoms of O for every 1 atom of N
1 mole of NO2 : 2 moles of O atoms to every 1 mole of N atoms
If we know or can determine the relative number of moles of each element in a compound, we can determine a formula for the compound.

18. Types of Formulas Empirical Formula
The formula of a compound that expresses the smallest whole number ratio of the atoms present.
Ionic formula are always empirical formula
Molecular Formula
The formula that states the actual number of each kind of atom found in one molecule of the compound.

19. Therefore, the empirical formula = CH (the lowest whole number ratio)
Formulas
Example: molecular formula for benzene is C6H6 (note that everything is divisible by 6)
Therefore, the empirical formula = CH (the lowest whole number ratio)

20. Formulas (continued) Molecular: Empirical:
Formulas for molecular compounds MIGHT be the same as empirical (lowest whole number ratio).
Molecular:
H2O
C6H12O6
C12H22O11
(Correct formula)
Empirical:
H2O
CH2O
C12H22O11
(Lowest whole number ratio)

21. To obtain an Empirical Formula
1. Determine the mass in grams of each element present, if necessary.
2. Calculate the number of moles of each element.
Divide each by the smallest number of moles to obtain the simplest whole number ratio.
If whole numbers are not obtained* in step 3), multiply through by the smallest number that will give all whole numbers
* Be careful! Do not round off numbers prematurely

22. Ex) A sample of a brown gas, a major air pollutant, is found to contain 2.34 g N and 5.34g O. Determine a formula for this substance.
Convert grams to moles
moles of N = 2.34g of N = moles of N
14.01 g/mol
moles of O = g = moles of O
16.00 g/mol
Formula:

23. Assume 100 g sample, so 38.67 g C x 1mol C = 3.22 mole C 12.0 g C
Calculate the empirical formula of a compound composed of % C, % H, and %N.
Assume 100 g sample, so
38.67 g C x 1mol C = 3.22 mole C g C
16.22 g H x 1mol H = mole H g H
45.11 g N x 1mol N = 3.22 mole N g N
Now divide each value by the smallest value

24. Example The ratio is 3.22 mol C = 1 mol C 3.22 mol N 1 mol N
The ratio is mol H = 5 mol H mol N mol N
= C1H5N1 which is = CH5N

25. Practice
A compound is % P and % O. What is the empirical formula?
Caffeine is 49.48% C, 5.15% H, 28.87% N and % O. What is its empirical formula?
= P2O5
= C4H5N2O

26. Empirical to molecular
1. add up the mass of the empirical formula.
2. Divide the mass of the molecular formula by the mass of the empirical formula.
3. Multiply the empirical formula by that number.

27. Calculation of the Molecular Formula
A compound has an empirical formula of NO2. The colorless liquid, used in rocket engines has a molar mass of 92.0 g/mole. What is the molecular formula of this substance?
Formula Mass
N = 14 g * 1 mol = 15 g/mol
O = 16 g * 2 mol = 32 g/mol
NO2 = 46 g/mol
= 92 g/mol
46 g/mol
= 2
Molecular Formula is N2O4

28. Molecular Formula Practice
The empirical formula for a compound is CH2O. Find the Molecular formula is the compounds Molecular Weight or Molar Mass is 180 g/mol?

29. End

30. Calculating with Moles-Stoichiometry

31. Chocolate Chip Cookies!!
1 cup butter
1/2 cup white sugar
1 cup packed brown sugar
1 teaspoon vanilla extract
2 eggs
2 1/2 cups all-purpose flour
1 teaspoon baking soda
1 teaspoon salt
2 cups semisweet chocolate chips
Makes 3 dozen
How many eggs are needed to make 3 dozen cookies?
How much butter is needed for the amount of chocolate chips used?
How many eggs would we need to make 9 dozen cookies?
How much brown sugar would I need if I had 1 ½ cups white sugar?

32. Stoichiometry
From the Greek words stoicheion meaning “element” and metron meaning “measure” measurement of the elements
Stoichiometry deals with the mass relationships between chemical elements

33. Mole Ratio
A mole ratio is a conversion factor that relates the amounts of moles in any two substances involved in a chemical reaction
Because any balanced chemical equation shows us the relative numbers of moles of the reactants and products we can write mole ratios for all chemical species given
Ex. 2HgO(s)  2 Hg (l) + O2(g) The equation tells us there are 2 moles of Mercury (II) Oxide, 2 moles of Mercury, and 1 mole of Oxygen gas

34. Chemistry Recipes
Looking at a reaction tells us how much of something you need to react with something else to get a product (like the cookie recipe)
Be sure you have a balanced reaction before you start!
Example: 2 Na + Cl2  2 NaCl
This reaction tells us that by mixing 2 moles of sodium with 1 mole of chlorine we will get 2 moles of sodium chloride
What if we wanted 4 moles of NaCl? 10 moles? 50 moles?

35. Mole Ratio
In Moleville…the mole ratio is needed to pass the guards to jump on the mole bridge to change from one substance to another
Always start with a Balanced equation!
Write what you are given with the label
Your mole ratio should be the ratio of the moles of what you want to what you are given.
Ex) Given 3 moles of H2 how many moles of water can be produced?
2H2 + O2 → 2H2O

36. Mole Ratios
Example: How many moles of chlorine are needed to react with 5 moles of sodium (without any sodium left over)?
2 Na + Cl2  2 NaCl
5 moles Na 1 mol Cl2
2 mol Na
= 2.5 moles Cl2

37. Mole-Mass Conversions
Most of the time in chemistry, the amounts are given in grams instead of moles
We still go through moles and use the mole ratio, but now we also use molar mass to get to grams
Example: How many grams of chlorine are required to react completely with 5.00 moles of sodium to produce sodium chloride?
2 Na + Cl2  2 NaCl
5.00 moles Na 1 mol Cl g Cl2
2 mol Na 1 mol Cl2
= 177g Cl2

38. Mass-Mole 10.0 g H2O 1 mol H2O 2 mol C2H6 18.0 g H2O 6 mol H20
We can also start with mass and convert to moles of product or another reactant
We use molar mass and the mole ratio to get to moles of the compound of interest
Calculate the number of moles of ethane (C2H6) needed to produce 10.0 g of water
2 C2H6 + 7 O2  4 CO2 + 6 H20
10.0 g H2O 1 mol H2O 2 mol C2H6
18.0 g H2O 6 mol H20
= mol C2H6

39. Mass-Mass Conversions
Most often we are given a starting mass and want to find out the mass of a product we will get (called theoretical yield) or how much of another reactant we need to completely react with it.
1. Change grams to moles
2. Do your mole ratio to switch substances
3. Change moles to grams of that substance

40. Mass-Mass Conversion 2.00g N2 1 mol N2 2 mol NH3 17.06g NH3
Ex. Calculate how many grams of ammonia are produced when you react 2.00g of nitrogen with excess hydrogen.
N2 + 3 H2  2 NH3
2.00g N mol N mol NH g NH3
28.02g N2 1 mol N mol NH3
= 2.4 g NH3

41. Practice
How many grams of calcium nitride are produced when 2.00 g of calcium reacts with an excess of nitrogen? 3Ca + N2 → Ca3N2

42. Limiting Reactants and Percent Yield

43. Limiting Reactants and Percent Yield
In the lab as well as industry, reactions are rarely carried out with exactly the required amount of each reactants. In most reactions, one or more reactants is in excess.

44. Limiting Reactants
The substance that is used up first is the limiting reactant.
The substance that is left over after the reaction is the excess reactant.
If you have 50 bicycle frames and 80 bicycle wheels, how many bicycles can be made. What limits the number of bicycles?

45. Example Problem P4 +O2 → P4O10
The reaction between solid white phosphorus and oxygen produces solid tetraphosphorus decoxide. Determine the mass of tetraphosphorus decoxide formed if 25.0 g of phosphorus (P4) and 50.0 g of oxygen are combined.
P4 +O2 → P4O10

46. When given masses of 2 reactants you need to determine the limiting reactant so go to moles of the product first. The one that makes the least moles is limiting.
Then change the moles to grams of the product to determine number of grams produced.

47. Percent Yield
Similar calculations are made to determine the success of a chemical reaction because most reactions do not succeed in producing the predicted amount of product.
Reactions fail to go to completion
Liquid adheres to containers or evaporates
Solids get stuck in filter paper or lost in the purification process.

48. Percent Yield
Theoretical yield- the maximum amount of product that can be produced from a given amount of reactant. (stoichiometry)
Actual yield-the amount of product actually produced when the reaction is carried out. (experimental)
Percent yield-the ratio of the actual yield to the theoretical yield as a percent.

49. Percent Yield
From the previous problem, the theoretical yield was calculated to be g P4O10. When the reaction was conducted in the lab 52.5 g of product was collected.
The percent yield is: