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Chemical Formulas and Molar Masses

Revisit old concepts and explore new ideas in determining molar mass, molecular weight, formula mass, and percentage composition. Practice calculating molecular formulas. Learn to determine empirical formulas from percentage composition and molecular weights. Includes examples and homework assignments.

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Chemical Formulas and Molar Masses

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  1. Chemical Formulas and Molar Masses A few old ideas revisited and a few new.

  2. Remember way back in Unit 2, we discussed the molar mass of elements • Back then, we discussed that the masses that appear on the periodic table are the Average atomic mass (atomic weight) of the isotopes of the elements. Remember this: Copper – 63 occurs naturally 72.7% of the time Copper – 65 occurs naturally 27.3% of the time How did we get the average atomic mass? 63 x 0.727 + 65 x 0.273 = 63.5 amu

  3. Well, we can use the masses on the periodic table to calculate the Molar Mass of various substances • Otherwise known as: • Molecular Weight • Molecular Mass • Formula mass • Formula weight • It is defined as the mass in grams of one mole of a substance

  4. How to determine Molar Mass (Molecular Weight) 1. Multiply the quantity of each element involved in the substance by the atomic mass (from the periodic table) of those elements 2. Add the masses calculated in step one together This is the molar mass of the substance in g/mol or amu

  5. Examples • Find the molecular weight of the following substances: • Ozone, O3 • A phosphate ion, PO4-3 • Arsenic trichloride • Ammonium carbonate, (NH4)2CO3

  6. Now practice these ones • Iron (III) Oxide, Fe2O3 • Magnesium Chloride, MgCl2 • Methane, CH4 • Water, H2O • Copper (II) acetate, Cu(C2H3O2)2 • Iron (II) Oxide

  7. Percentage Composition • The percent of each element that is in a compound • It gives us one way to check the identity of a newly created compounds • To determine percentage composition, divide the total mass of each element in a compound by the molecular weight of the compound and multiply by 100.

  8. Examples • What is the percentage composition of each element in the following? • Methane, CH4 • Water, H2O • Copper (II) acetate, Cu(C2H3O2)2 • A phosphate ion, PO4-3 • Lithium oxalate, Li2C2O4

  9. Empirical and Molecular Formulas(Going the other way) • Empirical Formula – the simplest ratio of the elements in a compound • For example, the simplest ratio of H2O2 is HO. • To determine the empirical formula, you will be given the percentage composition of the compound

  10. Example • Determine the empirical formula of a compound containing 69.94% Fe and 30.06% O. • Step 1 – Assume that you have 100 g total of the compound. Note, this is just an assumption that makes the math easier. Any mass of a compound could have been used to determine the percentages, but 100 g makes the math easier.

  11. Determine the empirical formula of a compound containing 69.94% Fe and 30.06% O. • Step 2 – Turn each mass into moles by dividing each elemental mass by its atomic mass. • Step 3 – Divide each mole value by the smallest mole value • Step 4 – If all of the numbers are not close to a whole number, multiply each by a whole number that will make each the smallest whole number possible

  12. Try These • Determine the empirical formula of a compound containing 30.9% Na, 47.7% Cl, and 21.5% O. • Determine the empirical formula of a compound containing 56.6% K, 8.7% C, and 34.7% O. • Determine the empirical formula of a compound containing 40.5% Zn, 19.9% S, and 39.6% O.

  13. Molecular Formulas • The true ratio of the elements in a compound. • To complete these type of problems, you will need the same information given to you for the empirical formula problems, but you will also be given the molecular weight of the compound

  14. Steps • Step 1 – Calculate the empirical formula (sometimes this will be given). • Step 2 – Calculate the molecular weight of the empirical formula. • Step 3 – Divide the molecular weight that was given by the calculated mass of the empirical formula. • Step 4 – Multiply each subscript of each element in the empirical formula by the resulting number from step 3

  15. Examples • Determine the molecular formula of each of the following compounds: • 85.6% C and 14.4% H with a molecular weight of 70.0 g/mol • 40.0% C, 6.7% H, and 53.3% O with a molecular weight of 60.0 g/mol • 34.3% Na, 17.9% C, and 47.8% O with a molecular weight of 134 g/mol

  16. Homework • Page 253 – 42, 44, 45, 47, 49, 50, 52, 53, 55, 57

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