1. Math in Chemistry

2. Percent Composition Purpose:
Can be used to figure out chemical formulas.

3. I. Percent Composition Two different types of problems:
1) Masses are given
2) No Masses are given

4. Masses are Given Steps to solve problem:
1) Add given masses to get total mass for one compound
2) Divide mass of each element by the total mass
3) Multiply by 100 to get the percent

5. Masses are Given Examples
1) A sample of Silver sulfide was found to contain 29.0 g of Ag and 4.30 g of S. Calculate the percent composition.

6. Masses are Given Examples
2) g of Na combines completely with 77.4 g of O. Calculate the percent composition.

7. No Masses Given Steps to solve problem:
1) Assume you have 1 mole of the compound and calculate its molar mass
2) Determine the molar masses of each element in the compound
3) Divide the molar mass for the element by the molar mass of the compound
4) Multiply by 100 to get the percent

8. No Masses Given Examples
1) Calculate the percent composition of C and H in ethane, C2H6.

9. No Masses Given Examples
2) Calculate the percent composition of sodium hydrogen sulfate.

10. II. Calculate the Mass of an Element in a Compound
Steps to solve the problem:
1) Find the molar mass of the compound
2) Find the percent composition of the element
3) Set up a conversion factor problem
Given mass (g compound) x percent composition/100 (g compound)

11. Calculate the Mass of an Element in a Compound Examples
1) Calculate the mass of hydrogen in the following:
A) 350 g C3H8
B) 20.2 g NaHCO3
C) 378 g HCN

12. III. Empirical Formulas
Empirical formulas are the lowest whole number ratio of the atoms of the elements in a compound.
Molecular Formula
C2O4 Na2O2
Empirical Formula
CO2 NaO

13. III. Empirical Formulas
Steps to solve a problem:
1) If the problem is given in percents, assume 100 g of the compound. This lets you easily convert the percents to grams (10% of 100g = 10g)
2) Convert the mass of each element to moles of that element
3) Find the smallest whole number ratio between the moles of the elements by dividing each molar mass by the smallest molar mass present.
4) The numbers you get as answers tell how many atoms of that element are present in the compound. If the numbers do not come out whole, round to the nearest whole number.

14. Empirical Formulas Examples
1) What is the empirical formula of a compound that is 49.6% nitrogen and 50.4% oxygen?

15. Empirical Formulas Examples
2) What is the empirical formula for a compound that is 79.8% C and 20.2% H?

16. Empirical Formulas Examples
3) What is the empirical formula of a compound that is 67.6% Hg, 10.8% S, and 21.6% O?

17. IV. Molecular Formulas
Usually the empirical formula is the molecular formula for a compound. When it is not, the molecular formula is defined as the elements and number of atoms that are contained in a compound.

18. IV. Molecular Formulas
Molecular formulas are always multiples of empirical formulas.
CH3 C2H6

19. IV. Molecular Formulas Steps to solve the problem:
1) Determine the empirical formula
2) Divide the molecular mass by the empirical formula mass to get a ratio.
3) Multiply the elements’ subscripts by the number you get in step 2.

20. Molecular Formulas Examples
1) Calculate the molecular formulas of the following compounds:
Molecular Mass Empirical Formula
60 g CH4N
78 g NaO
181.5 g C2HCl

21. Molecular Formulas Examples
2) The compound methyl butanoate smells like apples. Its percent composition is 58.8% C, 9.8% H, and 31.4% O. If the molecular mass is 102 g/mol, what is the molecular formula?

22. Molecular Formulas Examples
3) You find 7.36 g of a compound has decomposed to give 6.93 g of oxygen. The rest is hydrogen. If the molecular mass is 34.0 g/mol, what is the molecular formula?