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The Mole

The Mole. Counting Really, Really Large Numbers of Really, Really Small Things. Unit 5 Chapter 8. In The Beginning. In 1811, an Italian teacher & scientist, Amedeo Avogadro (1776-1856) Proposed that equal volumes of gases were composed of identical numbers of particles. Poor Amedeo.

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The Mole

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  1. The Mole Counting Really, Really Large Numbers of Really, Really Small Things. Unit 5 Chapter 8

  2. In The Beginning In 1811, an Italian teacher & scientist, Amedeo Avogadro (1776-1856) Proposed that equal volumes of gases were composed of identical numbers of particles.

  3. Poor Amedeo In his paper, he distinguished between atoms and molecules (though those words were used interchangeably at the time). His paper received little attention.

  4. The Rise of Avogadro In 1860, four years after his death, Avogadro’s theory was proven to be correct. Years later, in honor of his work, the number of particles calculated to be in an equivalent mass was designated as Avogadro’s number (NA). Avogadro is currently accepted as one of the founders of atomic theory.

  5. Enter The Mole In 1893, German chemist, Wilhelm Ostwald coined the term mole (Mol as an abbr. of the German word Moleküle) for the number of atoms in an equivalent mass of atoms or molecules. Scientists had been using equivalent mass for a long time.

  6. Current SI Definition 1. The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in… Exactly 12 grams of carbon-12; Its symbol is “mol”.

  7. Current SI Definition (cont’d) 2. When the mole is used, Elementary entities must be specified & may be: atoms, molecules, ions, electrons, other particles, or specified groups of such particles.

  8. This is Acceptable The current accepted value for the number of particles in a mole is 6.02214129(27)x1023. For simplicity, most people use 6.02x1023. (Use 6.022x1023 for 4 S.F.!) Written out in long form, this is: 602,000,000,000,000,000,000,000 602 with 21 zeroes after it 602 sextillion!

  9. Moles Everywhere This number is so famous (in academic and scientific circles) that October 23rd is the official “Mole Day” Sometimes, another “Mole Day” is instituted because 1st year chem classes usually do not learn about the mole in time – it’s June 2nd. There is even a website devoted to it: www.moleday.org Every year they have a different theme for Mole Day…this past year’s theme was:

  10. Everyday Uses In practice, it does not really matter how many particles are in a mole. What is important is our ability to use the mole to count equivalent amounts of atoms or molecules. Almost every calculation goes through the mole when you are comparing things in chemistry.

  11. Equivalent Units Recall that the numbers on the periodic table are in Daltons (or amu’s). Since Daltons are so small, they are impractical to use – you simply cannot weigh out 125 Da of anything. Our scales are not nearly sensitive enough.

  12. So That’s Why We Use a Scale Since the mass of 1 atom of Carbon-12 is defined as 12.00 Daltons and the mass of 1 mole of Carbon-12 atoms is defined as 12.00 grams, The masses on the Periodic Table are in two units – Daltons (Da) & grams per mole (g/mol). We typically use g/mol because we can weigh out amounts in grams.

  13. Apples & Oranges In order to compare amounts of elements, teachers like to say, “You cannot compare apples to oranges.” In chemistry, this means something like this: If you have 15 grams of carbon and 15 grams of silicon, do you have the same number of atoms of each?

  14. Applesauce…with Orange Peels The answer is no – there are more atoms in 15 grams of carbon than there are in 15 grams of silicon. The only way to know for sure is to compare “apples to apples.” This means we have to convert each substance into moles and then compare the moles.

  15. Calculating Moles The number of moles of a substance is abbreviated “n” Calculated by dividing the mass of substance by the weight on the PT (At Wt or Mn): n = m * 1 mol/Mn 15 g carbon * 1 mol C = 1.25 moles C 12.01 g C 15 g silicon * 1 mol Si = 0.53 moles Si 28.09 g Si

  16. Counting the Little Ones… If we want to find out how many atoms of carbon and silicon, we have to multiply by NA 1.25 moles C * 6.02x1023 = 7.53x1023 atoms 0.53 moles Si * 6.02x1023 = 3.19x1023 atoms There are 4.34x1023 more atoms in 15 g Carbon than in 15 g Silicon That’s 434,000,000,000,000,000,000,000 more!

  17. Interconverting X atoms * 1 mole NA X grams * 1 mole Mn Multiply by Mn Multiply by NA n moles * Mn 1 mole n moles * NA 1 mole Moles (n) Divide by Mn Divide by NA Mass (in grams) # of Particles

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