1. Unit 4: Moles, Chemical Composition, Stoichiometry
Chapter 7 & 9

2. Key Concepts Moles Avogadro’s Number Molar Mass Molar Volume
Percent Composition
Empirical Formula
Molecular Formula
Stoichiometry
Limiting Reactant, Excess Reactant
Theoretical Yield, Percent Yield

3. Moles Mole: SI base unit of measurement for the amount of a substance
 It provides a bridge between the atom and the macroscopic amounts of material that we work with in the laboratory.  It allows the chemist to weigh out amounts of two substances, say iron and sulfur, such that equal numbers of atoms of iron and sulfur are obtained. 
Dozen: 12 of something
Mole: 6.02x1023 of something (particles)
Moles and particles are units.
We can convert between these units by using this conversion factor.

4. Avogadro’s Number Carbon with 6P & 6N, atomic mass = 12amu
Avogadro massed 12g of carbon-12 and counted the atoms. He counted 6.02 x 1023 atoms and called it 1 mol.
Avocado kid vine
1 mol = x 1023 particles
(atoms/molecules)

5. Avogadro’s Number
Avogadro’s Number: one mole of a substance contains 6.02 x 1023 particles (atoms, molecules)
1 mol x 1023 particles
6.02 x 1023 particles mol

6. “It’s an avocado… Thanks.”
My student sent me this over summer…
“It’s an avocado… Thanks.”

7. Avogadro’s Number

8. Moles, Avogadro’s Number
How many moles are in 9.0 x 1021 atoms of carbon?
How many molecules are in 2.1 moles of carbon dioxide?
9.0 x 1021 atoms C x 1 mol/6.02 x 1023 atoms = mols C
2.1 mol CO2 x 6.02 x 1023 molecules/1 mol = 1.26 x 1024 molecules CO2

9. Moles, Avogadro’s Number
How many moles are in 4.35 x 1024 molecules of chlorine?
How many molecules are in 1.6 moles of HCl?
4.35 x 1024 molecules Cl x mol Cl = 7.23 mol Cl
6.02 x 1023 molecules Cl
1.6 mol HCl X x 1023 molecules HCl = 9.63 x 1023 HCl molecules
1 mol HCl

10. Molar Mass
Molar Mass: mass (g) of one mole (mol) of a substance (g/mol)
Atomic mass (g) = 1 mol
(add up atomic masses from the periodic table)
Atomic Mass (g) mol
1 mol Atomic Mass (g) Or

11. Molar Mass What is the mass of 0.0297 moles of NH3?
How many moles are in 3.5 grams of H2O?
mol NH3 x 17g/1mol = g NH3
3.5g H2O x 1mol/18g = mol H2O

12. Molar Mass What is the mass of 1.946 moles of NaCl?
How many moles are in 3.0 g of sodium hydroxide (NaOH)?
1.946 mol NaCl x g NaCl = g NaCl
1 mol NaCl
3.0 g NaOH x 1 mol NaOH = mol NaOH
40 g NaOH

13. Molar Volume
Molar Volume: volume (L) of one mole (mol) of a gas at STP (L/mol)
Volume 22.4L = 1 mol at STP (standard temperature & pressure)
22.4L Or mol
1 mol L
Standard Temperature: 0°C
Standard Pressure: 1atm

14. Molar Volume How many moles of O2 occupy 56L at STP?
How many liters will 0.25 mol of N2 occupy at STP?
56L O2 x 1 mol O2 = 2.5 mol O2
22.4L O2
0.25 mol N2 x 22.4L N2 = 5.6L N2
1 mol N2

15. Molar Volume How many moles does 5.6L of Cl2 occupy at STP?
How many liters does 2 moles of H2 occupy at STP?
5.6L Cl2 x 1 mol Cl2 = 0.25 mol Cl2
22.4L Cl2
2 mol H2 x 22.4L H2 = 44.8L H2
1 mol H2

16. Molar Volume Lab
Collect the hydrogen gas product in the top of the Eudiometer

17. 6.02 x 1023 particles 1 mol Atomic Mass (g) 22.4L
Conversion Factors
Avogadro
Molar Mass
Molar Volume
6.02 x 1023 particles 1 mol Atomic Mass (g) 22.4L
STP = standard temperature and pressure
standard temp is 0°C (freezing point of water)
standard pressure is 1 atm (sea level)
*Molar Volume is for gases at STP

18. Island Hopping Mass (g) Island Particle Volume (L)
(atoms/molecules)
Volume (L)
Convert between mass, volume, and particles.

19. Converting Between Mass (g), Volume (L), Particles (atoms or molecules)
What is the mass of 9L of He?
How many liters do 4.1 x 1023 molecules of F2 occupy?
How many molecules are there in 5g of carbon dioxide?
9L He x 1mol He x 4.0g He = 1.61g He
22.4L He 1mol He
(4.1x1023 atoms F2) x mol F x 22.4L F2 =
(6.02x1023 atoms F2) 1mol F2
5g CO2 x 1mol CO2 x (6.02x1023 molecules CO2) =
g CO mol CO2

20. This is how you get percent composition!
Analytical Chemistry
Analytical Chemistry: the chemical components of a substance are separated, identified, and quantified.
This is how you get percent composition!
Ex. coffee is made up of 50%C, 5%H, 16%O, and 29%N.
Qualitative Analysis: determines the presence or absence of a compound, but not the mass or concentration (quantity).
-uses chemical tests such as acid test for gold, flame test for metals, Kastle-Meyer test for blood.
Quantitative Analysis: gravimetric analysis determines the amount of material by weighing the sample before and after some transformation. volumetric analysis determines the concentration of an analyte using titration.
Instrumental Analysis: uses instruments, ex: spectroscopy (measures the molecules interaction with electromagnetic radiation), mass spectrometry (measures molecules mass-to-charge ratio using electric and magnetic fields), electrochemical analysis (measures potential (volts) and current (amps) in an electrochemical cell containing the analyte), thermal analysis (calorimetry, measures the interaction of a substance with heat), etc.

21. Percent Composition
Percent Composition: the percentage by mass of each element in a compound
% Oxygen = Mass of Oxygen x 100% = x 100% = 89%
Total Mass
% Hydrogen = Mass of Hydrogen x 100% = x 100% = 11%
Total Mass

22. Percent Composition
Find the percent composition of carbon and hydrogen in methane (CH4)
Molar Mass of CH4 =
% C =
% H4 =
Molar Mass of CH4 = (1.001 x 4) = g
% Carbon = mass of C/ mass of CH4 x 100% = / x 100% = 75%
% Hydrogen = mass H/ mass CH4 x 100% = 4.001/ x 100% = 25%

23. Empirical Formula Empirical Formula: simplified formula
The simplest whole-number ratio of atoms in a compound
Ex. AB2C is an empirical formula, but A3B6C3 is not
Use percent composition to calculate empirical formula
Convert to moles (%  g  mol)
Divide by the smallest number of moles

24. Empirical Formula
Ex. Chemical analysis of a substance identified its composition as 60.0% C, 13.4% H, and 26.6% O. Calculate the empirical formula.
C = 60.0% = 60g x 1mol/12g = 5mol 5mol/1.7mol = 3mol C
H = 13.4% = 13.4g x 1mol/1g = 13.4mol 13.4mol/1.7mol = 8mol H
O = 26.6% = 26.6g x 1mol/16g = 1.7mol 1.7mol/1.7mol = 1mol O
C3H8O

25. Molecular Formula Molecular Formula: actual formula
The actual number of atoms present, but they’re still in the same ratio as the empirical formula
Ex. The empirical formula of Glucose is CH2O, but its molecular formula is C6H12O6
Use empirical formula to calculate molecular formula
Calculate the number of times bigger the molecular formula is compared to the empirical formula
Experimental Molar Mass = # of times bigger
Molecular Molar Mass
Bellwork:
What is the empirical formula of 43.6% P and 56.4% O?
Experiment
Periodic Table

26. Molecular Formula
Ex. The empirical formula of a compound is P2O5, and its experimental molar mass is 284g/mol. What is its molecular formula?
Molecular Molar Mass = (30.97x2) + (15.999x5) = g/mol
Experimental Molar Mass = g/mol = 2 times bigger
Molecular Molar Mass g/mol
2(P2O5) = P4O10

27. Percent Composition  Empirical Formula  Molecular Formula
Ex. The percent composition of a molecule is 25.9% Fe and 74.1% Br with an experimental molar mass of 430g/mol. What is the empirical formula? What is the molecular formula?
Fe = 25.9g x 1mol/55.85g = 0.464mol 0.464mol/0.464mol= 1mol Fe
Br = 74.1g x 1mol/79.90g = 0.927mol 0.927mol/0.464mol= 2mol Br
FeBr2
Experimental Molar Mass = g/mol = 2 times bigger
Molecular Molar Mass g/mol
Fe2Br4

28. Percent Composition  Empirical Formula  Molecular Formula
Convert to moles (%  g  mol)
Divide by the smallest number of moles
Molecular Formula:
Experimental Molar Mass = # of times bigger
Molecular Molar Mass

29. Fruit Loops = Atoms
2 red fruit loops react with 3 blue fruit loops to make a molecule that contains 2 red and 3 blue fruit loops.
Write a balanced equation
Make your balanced equation with fruit loops
Each group gets a sandwich bag of fruit loops.
2R + 3B  1R2B3

30. Fruit Loops = Atoms
If you had 8 red, how many blues would it react with?
How many products could you make?
If you had 15 blue, how many reds would it react with?
8R + 12B  4R2B3
10R + 15B  5R2B3

31. Island Hopping Mass (g) Island Particle Volume (L)
(atoms/molecules)
Volume (L)
Convert between mass, volume, and particles.

32. Hopping Between Islands
Mol
Particles
(atoms/molecules
Volume (L)
Mass (g)
Island hopping for one kind of atom/molecule.
Use Avogadro’s conversion factor, molar mass conversion factor, and molar volume conversion factor to convert between particles, mass, and volume.

33. Hopping Between Islands
Mol
Particles
(atoms/molecules
Volume (L)
Mass (g)
Mol
Particles
(atoms/molecules
Volume (L)
Mass (g)
Use the mole ratio to convert between different atoms/molecules.

34. Hopping Between Islands
ATOM/MOLECULE
ATOM/MOLECULE
Mol
Particles
(atoms/molecules
Volume (L)
Mass (g)
Mol
Particles
(atoms/molecules
Volume (L)
Mass (g)
Avogadro
Avogadro
Molar Mass
Mole Ratio
Molar Mass
Molar Volume
Molar Volume

35. Stoichiometry
Stoichiometry: calculating the amount of reactants and products in chemical reactions
Mole Ratio: the coefficients in a chemical equation represent its ratio of moles
Ex. N H2  2NH3
1mol of N2 reacts with 3mol of H2, which yields 2mol of NH3
This balanced equation will always be a 1:3:2 ratio.
*The mole ratio for a reaction is always the same!

36. Stoichiometry
N H2  2NH3
Ex. What volume of H2 gas would you need to make 8.5g of NH3?
8.5g NH3 X 1mol NH3 x 3mol H2 x 22.4L H2 = 16.8L H2
17g NH3 2mol NH3 1mol H2

37. Limiting Reactant
Limiting Reactant: the reactant that gets used up first; it limits the amount of product that can form
Theoretical Yield: produced by limiting reactant
Excess Reactant: the reactant that is not completely used up in a reaction
Example: boys + girls  couples
Less girls = limiting reactant [makes theoretical yield]
More boys = excess reactant
The limiting reactant (girls) determines how many couples can be made. The excess reactant (boys) is leftover.

38. Limiting Reactant Convert each reactant to g of product
The reactant that makes the least amount of product is the LR, the other reactant is the ER
The product made with the LR is the TY

39. Limiting Reactant
Ex. If you combined 56g of iron and 32g of oxygen, how much iron (III) oxide can be made?
4Fe + 3O2  2Fe2O3
Fe: 56g Fe x 1mol Fe x 2mol Fe2O3 x g Fe2O3 = 80.1g Fe2O3
55.85g Fe mol Fe mol Fe2O3
O2: 32g O2 x 1mol O2 x 2mol Fe2O3 x g Fe2O3 = 106g Fe2O3
32g O mol O mol Fe2O3

40. % Yield = Actual Yield x 100%
Percent Yield
Percent Yield: tells you how well an experiment went
% Yield = Actual Yield x 100%
Theoretical Yield
Given
Calculated
Ex. 90% yield means you got 90% of the product you were supposed to get

41. Percent Yield
Ex. You reacted 11.5g of sodium with excess chlorine gas, and it produced 13g of sodium chloride. What is the % yield?
Balanced chemical equation:
Theoretical yield:
% yield:
2Na + Cl2  2NaCl
11.5g Na x 1mol Na x 2mol NaCl x 58.5g NaCl = g NaCl
23g Na mol Na mol NaCl
% yield: 13g NaCl x 100% = 44.4%
29.25g NaCl