1. Chapter 3 Mathematics of Formulas and Equations
PAGES (46-56)

2. Aim: To Apply Simple Math Relations to the Solving of Chemical Problems

3. Gram Formula Mass Compounds are made up of more than one element
Gram formula mass- sum of the atomic masses of all of the atoms in the compound
To find gram formula mass (GFM):
List the elements present
Count how many of each element
Find each element’s atomic mass
Multiple the atomic mass by the number of each element
Add it together

4. Examples: Find each of the GFM
NaBr
H2O
Ca(OH)2
Mg(NO3) 2
(NH4)3PO4

5. Open your book to p. 47 (1-13) 4
4 = 254g/mol 3 4 1 2

6. Percent Composition
Composition of a substance within a compound, expressed as a percent
Table T
To find the percent composition:
Find the GFM of the compound(s)
Find the mass of the part in the question
Plug each part into the formula
Hydrates- Compound plus a water part
CuSO4 * 5H2O
Many times, they will ask for how much water there is within the hydrate

7. Examples: What is the percent composition in each compound?
Hydrogen in H2O
Carbon in CO2
Oxygen in Ca(OH)2
Nitrogen in Fe(NO3)3
Chlorine in KClO3

8. Examples: Find the amount of water within each hydrate
Na2CO3 * 10H2O
CaSO4 * 2H2O
CH2O * 3H2O
CuSO4 * 5H2O
Mg3(PO4) * 4H2O

9. Open your book to p (14-26)

10. The MOLE A collective noun Number of atoms present 6.02 x 1023
Dozen, gross, pair, etc.
Number of atoms present
Actually, it’s the number of atoms present in 12 grams of Carbon-12
6.02 x 1023
Avagadro’s number
GFM = 1 mol of a substance

11. Converting Between Grams and Moles
Set up a proportion:
1 mol = # of moles
GFM # of grams
Converting Grams to Moles:
Moles = number of grams x 1/GFM
Converting Moles to Grams:
Grams = number of moles x GFM/1

12. Examples: Convert the following grams to moles
Converting Grams to Moles:
Moles = number of grams/GFM
36g of H2O
185g of KCl
80.0g of C2H5Cl

13. Examples: Convert the following moles to grams
Converting Moles to Grams:
Grams = number of moles x GFM
3mol of CO2
2.1mol of NaCl
0.12mol of C2H6
1mol of H2O

14. Finding Molecular Formulas from Empirical Formulas
When the molecular mass and the empirical formula are known, you can find the molecular formula.
If the empirical formula is CH2, then you know that the empirical formula’s mass must be 14g.
If the molecular mass is 56g, it must be a multiple of the empirical.
So, divide the given mass by the GFM of the empirical formula
56 / 14 =
The answer will ALWAYS be a whole number.
Now, multiply each element by the number to find the molecular formula.
What is it?

15. Examples: Find the Molecular Formula
Molecular Mass = 36g Empirical Formula = H2O
Molecular Mass = 140g Empirical Formula = CH2
Molecular Mass = 180g Empirical Formula = CH2O
Molecular Mass = 60g Empirical Formula = NO

16. Mole Relations in Balanced Equations
The coefficients in a balanced equation equal the number of moles of each element/compound.
So, the ratios of reactants and products will ALWAYS stay the same.
For example,
2H2 + O2 => 2H2O
For every 2 moles of H2 used, there are 2 moles of H2O produced.
For every 2 moles of H2 used, there is 1 mole of O2 also used.
Now, if 4 moles of H2 used, how many moles of H2O are produced?
Remember that it’s the SAME ratio!
Also, if 3 moles of H2 used, how many moles of O2 are used?

17. Examples: CH4 + 2O2 => CO2 + 2H2O 2CO + O2 => 2CO2
What is the mole ratio of CH4 to O2?
So, how many moles of O2 are need for 3mol of CH4?
2CO + O2 => 2CO2
What is the mole ratio of CO to CO2?
So, how many moles of CO are needed to create 1mol of CO2?
Mg + 2HCl => MgCl2 + H2
What is the mole ratio of HCl to MgCl2?
So, if 3mol HCl were used, how many MgCl2 were created?