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Chemical Measurements and Mole Conversions: Understanding The Mole and Empirical/Molecular Formulas

Learn about the concept of the mole, its role in chemical measurements, and how to convert between moles, grams, and particles. Understand how to determine empirical and molecular formulas through percentage composition and ratio calculations.

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Chemical Measurements and Mole Conversions: Understanding The Mole and Empirical/Molecular Formulas

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  1. Unit 9 part 1: The Mole 9.1.1 Chemical Measurements 9.1.2 Mole Conversions 9.1.3 Empirical & Molecular Formulas

  2. 9.1.1 Chemical Measurements • Remember… way back in Ch 3… we talked about ATOMIC MASS- the total # of protons and neutrons in nucleus of the atom. The units for atom mass are amu. • Ex: Carbon – 12.011 amu This is also know as the atomic weight of the atom.

  3. Formula Mass If it’s not just an element, but a compound… then you have: • Formula Mass- the sum of the atomic masses of all the atoms in the compound. • Ex. CO2 C- 12.011 O- 15.999 X 2 44.009 amu – formula mass YOU TRY! • H2O • CH4Cl2

  4. Relating amu and moles We have a problem… how do you measure amus? How do you count atoms? They are so small… that you don’t. Instead, we use a number called the mole to help us. • A mole (of any element) – the number of atoms equal to the number of atoms in exactly 12.0 g of carbon-12.

  5. How does it work? • There is an actual # … it is 6.02 x 1023 atoms/gram. • Notice the amu- gram link? “The mass in grams of 1 mole of a substance is numerically equal to its atomic mass or formula mass in atomic mass units.” • If you have 1.0 g of H.. you have 6.02 x 1023 atoms of H • If you have 16.0 g O, you have 6.02 x 1023 atoms of O How many grams of Pb would you need in order to have 6.02 x 1023 atoms of Pb?

  6. Mole Construction • The mole can be used to count atomsin elements: • To count molecules in covalent bonds: • To count formula units inionic bonds:

  7. Avogadro’s Number • The Mole is also known as Avogadro’s Number (N) in honor of Amadeo Avogadro, the Italian chemist. How Big is Avogadro’s number? 602,000,000,000,000,000,000,000! if it was a stack of paper… it would reach beyond our solar system! …basketballs… it could create a planet the size of the earth! …grains of rice… would cover the earth to a depth of 246 feet!

  8. Molar Mass (links amu and mole) • “The mass in grams of 1 mole of a substance..” (M). • Molar Mass of… • Ca - 40 amu –> 40 g/mol • C - 12 amu -> 12 g/mol • H2O2 -34 amu -> 34 g/mol • NaCl - • CaCl2 - • C6H12O6 -

  9. 9.1.2 Gram-mole-Particle Conversions We can use the molar mass to convert moles to grams or grams to moles. Likewise, moles can be converted to # of particles via Avogadro’s #.

  10. Solve: • How many moles of NaCl equals 60g of NaCl? • How many grams of Zn(NO3)2 equals 5.5 mols? • How many atoms of Ag equals 6.2 mols? • How many mols of CaCO3 equals 5.36 x 1023 formula units? • How many grams equals 2.5 x 1023 formula units of NaCl?

  11. You can convert moles to volume (L). Use the ratio 22.4L = 1 mole. No matter what the gas is… it’s always the same ratio. • How many moles equals 65.5L of O2? • How many Liters of CO2 equals 10.3 mols? • How many L equals 42.0g H2O?

  12. 9.1.3Empirical & Molecular Formulas • We all know water has the molecular formula of H2O… but have you ever wondered how they figured that out? • Well… first we must understand PERCENTs.

  13. Percentage Composition • If you spend 6 hours at school every day for 20 days in a month. WHAT IS THE PERCENT OF TIME you spent at school for the whole month if there are 31 days in the month? • % = (actual time ÷ total time) x 100 • ((6x20) / (24 x 31)) X 100 = 16.13%

  14. What does this have to do with formulas? Similarly, we can calculate the percents of element mass in a compound. • “The mass of each element in a compound compared to the entire mass of the compound and multiplied by 100 percent is called the percentage composition of the compound.” • %= (mass of element÷ total mass of the compound) x 100

  15. For example… • If you have one mole of water (total mass of 18g), what is the percent composition of H and O? • Hint: You need to know the formula to solve for % comp if you don’t know the mass for each element. H2O- which means you have 2 mol H and 1 mol O. • Well… what is the mass of 2 mole H and what is the mass of 1 mol of O? • H – 2g & O- 16g • Now… calculate… (mass of element÷ total mass) x 100 • % H = (2/18)x 100 = 11% • % O = (16/18) x 100 = 89%

  16. Determining Empirical Formula • “A formula that gives the simplest whole-number ratio of the elements is called an empirical formula.” • We can use the ratio of the masses to determine this formula. • But… what’s a ratio? Think of the Domino’s commercial… people prefer Domino’s toasted sandwiches to Subway 2:1.

  17. If we know we have 11% H and 89% O what would the EMPIRICAL formula for water be? 1. Pretend you have a 100g sample. (change % to g) 2. Convert grams to moles. (molar mass) 3. Divide each mole by the SMALLEST mole #. 4. Round to a whole number and write the empirical formula H: 11% of 100g = 11g O: 89% of 100g = 89g H: 11g/1g = 11 mol H O: 89g/16g = 5.56 mol O H: 11/5.56 = 1.977 O: 5.56/5.56 = 1 H: 2 O: 1 H2O Working the H2O example backwards

  18. Molecular Formula • “The formula that gives the actual number of atoms of each element in a molecular compound is called the molecular formula.” • This will be the EMPIRICAL formula multiplied by a whole number. • Ex. HO x 2 = H2O2… hydrogen peroxide.

  19. How do you know what whole # to multiply the empirical formula by? • The question will always tell you the MOLAR MASS of the molecular compound. • Take the molar mass of the compound and divide by the empirical formula mass. Molar mass ÷ Empirical formula mass • The answer is the WHOLE # to multiply the empirical formula by.

  20. EX: Ribose has a molar mass of 150g/mol and a chemical composition of 40.0% C, 6.67% H and 53.3% O. What is the molecular formula? • First, find the empirical formula. 1. Pretend you have a 100g sample. 2. Convert grams to moles 3. Divide each mole by the SMALLEST mole #. 4. Round to a whole number and write the empirical formula • Now calculate the formula mass of the empirical formula. • Divide molar mass by formula mass. • Molar mass ÷ Empirical formula mass • Multiply the whole # by the empirical formula = molecular formula.

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