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Chapter 6

Chapter 6. Counting Atoms by Weighing The Mole. What is Chapter 6 all about?. Chapter 6 is all about keeping track of atoms. We are going to find out how to tell how many atoms there are in a certain amount of material.

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Chapter 6

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  1. Chapter 6 Counting Atoms by Weighing The Mole

  2. What is Chapter 6 all about? Chapter 6 is all about keeping track of atoms. We are going to find out how to tell how many atoms there are in a certain amount of material. We are going to figure out how much mass a certain amount of atoms will have. We will also learn a little bit about how we know the formulas of compounds.

  3. How can we keep track of atoms? They are too small to see, so we can’t count individual atoms. Also, they have extremely small masses. For example, the mass of a single carbon atom is 1.99 x 10-23 g.

  4. How can we keep track of atoms? To avoid using terms like 10-23 when describing the mass of an atom, scientists have created an easier unit to use. Atomic mass unit (amu) = 1.66 x 10-24g.

  5. How can we keep track of atoms? There are different masses for different isotopes of an element. The isotopes of carbon are 12C, 13C, 14C. And any sample of carbon has a mixture of these isotopes (always in the same proportion). Since each isotope has a different mass, we need to use an average mass for the carbon atoms. Average atomic mass - the weighted average of the masses of all the isotopes of an element.

  6. Average Atomic Mass Find the average atomic mass for the following: • Cu • H • C • O

  7. Using Mass to Count Atoms 1 carbon atom = 12.01 amu How many carbon atoms are in a 3.00 x 1020 amu sample?

  8. Using Mass to Count Atoms How many oxygen atoms are in a 3.00 x 1020 amu sample? How many nitrogen atoms are in a 3.00 x 1020 amu sample?

  9. Finding Mass from the Number of Atoms Calculate the mass, in amu, of a sample of aluminum that contains 75 atoms.

  10. Finding Mass from the Number of Atoms Calculate the mass, in amu, of a sample of calcium that contains 100 atoms. …a sample of 250 phosphorus atoms.

  11. What did we do?Where are we going? We have just learned how to use amu for counting atoms. However, in the lab, we need a much larger unit (grams) We are going to learn how to count atoms in samples with masses given in grams.

  12. The Mole A mole (abbreviated mol) is just a number (like dozen). Also called “Avogadro’s number.” 6.022 x 1023 One mole of something consists of 6.022 x 1023 units of that substance. A dozen eggs = 12 eggs A mole of eggs = 6.022 x 1023 eggs.

  13. The Mole Why do we have moles? 1 mole of an element is equal to its amu in grams. EX: Fluorine has a mass of 19.00 amu 1 mole of Fluorine = 19.00 g

  14. The Mole Find the mass (in grams) for a mole of each element: • Hydrogen • Lithium • Phosphorus • Sodium

  15. The Mole Find the number of moles in given amount of substance: • 10.81 g of boron • 10 g of carbon • 32 g of oxygen

  16. The Mole Find the number of atoms in a given amount of substance: • 10.81 g of boron • 10 g of carbon • 32 g of oxygen

  17. The Mole Find the mass (in grams) for the given amount of substance: • 2.13 x 1024 atoms Au • 3.14 x 1023 atoms C • 2.74 x 1023 atoms Fe

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