1 / 9

Empirical Formulas

Empirical Formulas. Chapter 6, Section 3. What is an Empirical Formula?. Up to this point, we have been using molecular formulas to describe the chemical composition of compounds. A molecular formula is the actual formula of a compound

mullisb
Download Presentation

Empirical Formulas

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Empirical Formulas Chapter 6, Section 3

  2. What is an Empirical Formula? • Up to this point, we have been using molecular formulas to describe the chemical composition of compounds. • A molecular formula is the actual formula of a compound • By contrast, an empirical formula is a formula of a compound expressing the smallest whole-number ration of atoms in a compound

  3. Examples of Empirical Formulas • Molecular Formula: C6H12O6 • Empirical Formula: CH2O • Molecular Formula: C6H6 • Empirical Formula: CH • Molecular Formula: C12H4Cl4O2 • Empirical Formula: C6H2Cl2O

  4. Calculations of Empirical Formulas • Remember, chemical formulas are a ratio of atoms in a molecule, but it is also a ratio of moles of atoms. • Therefore, you can calculate the empirical formula for a compound given the percent composition in the compound or mass in grams

  5. Steps for Calculating Empirical Formulas • If you are given percentages change the percentages into grams. If you start in grams, go to step 2 • Convert grams of each element to moles • Divide each mole value by the smallest number. If close enough, round to the nearest whole number. If you get decimals, see step 4 • Multiply your fractions by the smallest whole number that will make all you numbers whole numbers

  6. Example 1 • Calculate the empirical formula of a compound composed of 38.67% C, 16.22% H, and 45.11% N • C = 38.67 g, H = 45.11 g, N = 45.11 g • Mole Calculations • 38.67 g C x 1 mol C = 3.220 mole C 12.01 g C • 16.22 g H x 1 mol H = 16.09 mole H 1.01 g H • 45.11 g N x 1 mol N = 3.219 mole N 14.01 g N

  7. Example 1 cont • C is the smallest. Divide all number by 3.220 mol C • 3.220 mol C / 3.220 mol C = 1C • 3.219 mol N / 3.220 mol C = 1N • 16.09 mol H/ 3.220 mol C = 5 H Our empirical formal is CH5N

  8. Example 2 • You have a sample that contains 4.151 g of Al and 3.692 g of oxygen. Calculate the empirical formula. • Already done • Mole Calculations • 4.151 g Al x 1 mol Al = 0.1539 mol Al 26.98 g • 3.692 g O x 1 mol O = 0.2308 mol O 16.00 g

  9. Example 2 cont • Al is the smallest. Divide all by 0.1539 mol Al. • 0.1539 mol Al/ 0.1529 mol Al = 1 Al • 0.2308 mol O/ 0.1529 mol Al = 1.5 O • Multiplying by 2 will get you to whole numbers • 1 Al x 2 = 2 Al • 1.5 O x 2 = 3 O Empirical formula is Al2O3

More Related