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Unit 7 chemical Calculations

Unit 7 chemical Calculations. Ch 9: P 289 – 290 Ch 10: P 325 – 335 Ch 15: P 498. Percent composition. % composition determines how much each elements’ mass contributes to the compound mass. What is the % composition for K 2 CrO 4 ?. K: 39.1 (2) = 78.2 Cr: 52.00 O: 16(4) = 64

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Unit 7 chemical Calculations

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  1. Unit 7 chemical Calculations Ch 9: P 289 – 290 Ch 10: P 325 – 335 Ch 15: P 498

  2. Percent composition • % composition determines how much each elements’ mass contributes to the compound mass.

  3. What is the % composition for K2CrO4? • K: 39.1 (2) = 78.2 • Cr: 52.00 • O: 16(4) = 64 52.00 + 64.00 + 78.2 = 194.2 (total mass) K: (78.2/194.2) X 100 = 40.3 % Cr: (52/194.2) x 100 = 27.8 % O: (64/194.2) x 100 = 33.0 %

  4. Law of Definite Proportions Says that chemical compounds always contain the same proportions of elements by mass Ex: Water (H2O) will always be made of 2.0 g/mol of H and 16.0 g/mol of O. H = 1.0 g/mol O = 16.0 g/mol In water there are 2 H so 2(1.0 g/mol) = 2.0 g/mol and 1O = 16.0 g/mol

  5. Law of Multiple Proportions • Says that when elements combine, they do so in the ratio of small whole numbers Ex: CO2 can exist (2 is a whole number) CO1.5 can not exist

  6. Hydrates: • A hydrate is something that contains water • Hydrated compounds contain water among other elements • Helpful equation for finding the % of water:

  7. Examples: calcium chromate trihydrate is written as:CaCrO4 · 3H20 • The following masses are found on the periodic table: Ca = 40.1 H = 1.0 (2) =2.0 Cr = 52.0 O = 16.0 O= 16.0 (4) = 64.0 18.0 (3) 156.1 g/mol 54.0 g/mol

  8. Try the following on your own • copper(II)nitrite nonahydrate • CuSO4· 5H2O • CaCl2 · 6H2O

  9. Empirical Formula: • An empirical formula shows the smallest whole number ratio of the atom in the compound.

  10. Steps for solving an empirical formula A Use the % Composition B Assume all percentages are out of a 100g sample C Convert all grams into moles D Write formulas using the moles you found in c as subscripts E Divide all subscripts by the smallest mole amount F Make sure all numbers are small whole number ratios. If not, multiply so that they are.

  11. Examples: The % composition of a compound is 32.14% Al and 67.86% F. Calculate its empirical formula Lets use the steps from above: • % compositions is already given • 32.14% = 32.14g Al • 67.86% = 67.86g F .

  12. Al1.19F3.57 • Al F =AlF3 • AlF3 is already in whole numbers so we are done

  13. Now take the Empirical formula and determine the molecular formula: • The molecular formula can be the same as the empirical formula or a whole # multiple of its empirical formula Ex: CH =empirical CH2O = empirical • C2H2 = molecular C2H4O2 = molecular • C6H6 = molecular C6H12O6 = molecular

  14. Steps for finding molecular formula: a. Determine the empirical formula mass (empirical formula will be given or you will have to determine it) b. Divide the given molar mass by the mass of the empirical formula c. Multiply the subscripts in the empirical formula by the answer you got in step (b)

  15. Example: Empirical formula: CH3O. If the molar mass of the molecular mass is 62.0 g/mol, what is the molecular formula? • C – 12.0 • H – 1.0 (3) = 3.0 • O – 16.0 ________ • 31.0 g – empirical mass

  16. Molar mass/empirical mass • Take the empirical formula and multiply all the subscripts by 2 • CH3O (2) = C2H6O2

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