Unit 10 – The Mole

1. Unit 10 – The Mole • Essential Questions: • What is the relationship between a mole of a substance and its mass? • How can the mole of a substance be calculated? • How can the percent composition of a compound be determined? • How does the molecular formula of a compound compare with the empirical formula?

2. Molecular and Formula Masses • The sum of the masses of all the atoms in a compound • Molecular Mass – mass in a molecule (covalent) • Formula Mass – mass in a formula unit (ionic) • Unit is amu for either of them • Example: Formula Mass of CaCO3 • 1 atom of Ca = 40.08 amu = 40.08 amu • 1 atom of C = 12.01 amu = 12.01 amu • 3 atoms of O = 3 x 16.00 amu = 48.00 amu 100.08 amu

3. Example: Find the formula mass of (NH4)2SO3 2 N = 2(14.01 amu) = 28.02 amu 8 H = 8(1.01 amu) = 8.08 amu 1 S = 32.07 amu = 32.07 amu 3 O = 3(16.00 amu) = 48.00 amu 1 formula unit = 116.17 amu

4. Try these problems: • HNO3 • C6H10O5 • Al3(PO4)2 = 63.02 amu = 162.16 amu = 270.88 amu

5. Mole • A counting number (like a dozen) • 6.02 X 1023 (in scientific notation) • This number is named in honor of Amedeo Avogadro (1776 – 1856) • Discovered that no matter what the gas was, there were the same number of molecules present in the same volume

6. Mole – 6.02 x 1023 particles = 6.02 x 1023 C atoms = 6.02 x 1023H2O molecules = 6.02 x 1023NaCl formula units 6.02 x 1023 Na+ ions and 6.02 x 1023Cl– ions 1 mole C 1 mole H2O 1 mole NaCl

7. Avogadro’s Number as Conversion Factor Particles = Moles 6.02 x 1023 particles 1 mole Or Moles = Particles 1 mole 6.02 x 1023 particles Note that a particle could be an atom OR a molecule! You MUSTuse dimensional analysis for conversions! X X

8. Examples: How many molecules are in 3.5 moles of H2O? How many moles are present in 465 molecules of NO2? How many atoms of nitrogen are in 3.15 moles of NH3? How many atoms of chlorine are in .862 moles of MgCl2?

9. Molar Mass • Molar Mass- the mass of one mole of a substance • Unit is grams/mole (g/mole or g/mol) • Equivalent to the molecular mass in amu • Ex: molar mass of Iron = 55.85 g /mole molecular mass of Iron = 55.85 amu

10. Mass and Mole Relationships • Find the number of moles present in 56.7 g of HNO3. • Find the number of grams present in 4.5 moles of C6H10O5. • Find the number of moles present in 12.31 g of H2SO4.

11. Percent Composition • Finding what percent of the total weight of a compound is made up of a particular element • Formula for calculating % composition: Total amu of the element in the compound Total formula amu X 100%

12. Example Calculate the % composition of BeO

13. Example Calculate the % composition of Ca(OH)2

14. Example Calculate the % composition of Al(NO3)3

15. Chemical Formulas • 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). • 1 molecule NO2 : 2 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

16. Law of Multiple Proportions • When any two elements, A and B, combine to form more than one compound, the different masses of B that unite with a fixed mass of A bear a small whole-number ratio to each other • Example: • In H2O, the proportion of H:O = 2:16 or 1:8 • In H2O2, H:O is 2:32 or 1:16

17. Empirical vs. Molecular Formula • Empirical Formula - The formula of a compound that expresses the smallest whole number ratioof the atoms present.Ionic formulas are always empirical formulas • Molecular Formula - The formula that states the actual number of each kind of atom found in one moleculeof the compound.

18. Determine the Empirical Formula From the Molecular Formula All you need to do is reduce!! • C6H6 • Fe3(CO)9 • BaCl2 • P4O10

19. Determine the Molecular Formula from the Empirical Formula • Calculate the molar mass of the Empirical Formula. • Divide the molar mass of the Molecular Formula by the molar mass of the Empirical Formula • Multiply the numbers of each type of atom by that number

20. Determine the Molecular Formula from the Empirical Formula • Examples: • Molecular Mass: 26.04 g/mol • Empirical Formula: CH • Molecular Formula: C2H2 • Molecular Mass: 380.88 g/mol • Empirical Formula: SeO3 • Molecular Formula Se3O9

21. To Obtain Empirical Formula • 1. Assume the percent is out of 100 grams. That means you can change the % sign to grams. • 2. Calculate the number of molesof each element. • 3. 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 **Remember this** Percent to massMass to moleDivide by smallMultiply 'til whole

22. Calculating the Empirical Formula Example #1 • Given that a compound is composed of 60.0% Mg and 40.0% O, find the empirical formula.

23. Calculating the Empirical Formula Example #2 A compound is analyzed and is found to contain 13.5g of calcium, 10.8g of oxygen, and 0.675g of hydrogen. Calculate the empirical formula of this compound.

24. Calculating the Empirical Formula Example #3: NutraSweet is a zero calorie sweetener used in many food products. A sample is analyzed and it’s percent composition is as follows; 57.14% carbon, 6.16% hydrogen, 9.52% nitrogen, and the rest is oxygen. Calculate the empirical formula of NutraSweet.

25. Try this! A compound is found to contain 68.5% carbon, 8.63% hydrogen, and 22.8% oxygen. The molecular weight of this compound is known to be approximately 140.00 g/mol. Find the empirical and molecular formulas.