1. Chemistry Calculations
Moles
Notes #4b

2. Review Mole-mass Mass-mole Mole-particles Particles-moles
Particles – mass
Mass - Particles

3. Mole-Mole Calculations
Chemical Formulas and
Chemical Reactions
contain mole ratios

4. Mole-Mole Calculations
N2 + 3 H NH3
The most important part of this equation is that it tells us that 1 mol of nitrogen reacts with 3 mol of hydrogen to form 2 mol of ammonia.
The coefficients from the balanced equation are used to write conversion factors called mole ratios.

5. Mole Ratio
Mole ratios make a ratio with the number of moles of one element/compound to the number of moles of another element/compound.
N2 + 3 H NH3
1 mol N2
2 mol NH3
3 mol H2
2 mol NH3
2 mol NH3
1 mol N2
What are the other three mole ratios that can be made?

6. SAMPLE PROBLEM N2 (g) + 3H2 (g) 2NH3 (g)
How many moles of ammonia are produced when 0.60 mol of nitrogen reacts with excess hydrogen?
N2 (g) + 3H2 (g) NH3 (g) #1
2 mol NH3
1 mol N2
0.60 mol N2 1
= 1.2 mol NH3 x

7. SAMPLE PROBLEM For the reaction 4 Al (s) + 3 O2 (g) 2 Al2O3 (s)
Write the six mole ratios that can be derived from this equation.
How many moles of aluminum are needed to form 3.7 mol Al2O3?

8. Review Molar Mass
For an element
For a Compound

9. Mass-Mass Calculations
A sample is measured in grams and you are asked to find the number of grams of another substance in the reaction.
Uses molar mass and mole ratios to figure out the grams of the unknown.

10. SAMPLE PROBLEM N2 (g) + 3H2 (g) 2NH3 (g) 30.4 g NH3 5.40 g H2 1
Calculate the number of grams of NH3 produced by the reaction of 5.40 g of hydrogen with an excess of nitrogen.
N2 (g) + 3H2 (g) NH3 (g)
Want to know = g NH3
5.40 g H2 1
1 mol H2
2.02 g H2
2 mol NH3
3 mol H2
17.04 g NH3
1 mol NH3 x x x =
30.4 g NH3

11. Zn(s) + 2HCl(aq) ZnCl2 (aq) + H2 (aq)
SAMPLE PROBLEM
Zn(s) + 2HCl(aq) ZnCl2 (aq) + H2 (aq)
If 20.0 g of zinc react with excess hydrochloric acid, how many grams of zinc chloride can be produced?
41.7 g ZnCl2

12. 11.4 Empirical and Molecular Formulas / %Composition

13. Empirical and Molecular Formulas
% Composition= the % by mass of each element in a compound
Mass of element X 100= % by mass
Mass of compound

14. %Composition

15. Empirical Formula= the formula with the smallest whole number mole ratio of the elements in a compound
May or may not be the same as the molecular formula

16. Problem:
Given 3.08 % H, 31.61% P, and % O find the empirical formula (assume 100g of compound = 100%)

17. 3.05 mol H = 2.99 = 3
1.02
1.02 mol P = 1
4.08 mol O = 4
H3PO4

18. Empirical Formulas Using % Composition and assuming % are out of 100g
Convert % to grams and then to moles
Determine simplest whole number ratio of the elements in the compound by dividing by the lowest number of moles.

19. Empirical Formulas Continued
A compound contains 59.95% O and 40.05% S, determine the empirical formula for the compound.

20. Empirical Formulas Continued
Once the number of moles is determined, divide each mole quantity by the least number of moles.
3.74 mol O/1.25mol S = 2.99 mol O = 3 mol O
1.25 mol S/1.25 mol S = 1.00 mol S = 1 mol S
The empirical formula would be SO3

21. Molecular Formulas
Molecular formula specifies the actual number of atoms in a compound and is typically a multiple of the empirical formula, but may be the same as the emp. Formula.

22. Molecular formula CH2O = empirical formula C6H12O6 = molecular formula
(CH2O)n = (CH2O)6 = C6H12O6
Molecular formula = (empirical formula )n

23. Example
A compound has the empirical formula CH, its molecular formula mass is 26.04, determine the compounds molecular formula
26.04/13.02=2
Therefore the molecular formula is C2H2

24. Hydrates
We will perform a Hydrate lab to understand hydrates