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Chapter 10

Chapter 10. The Mole. Section 10-1. Atomic Mass – The mass of an atom relative to the mass of C-12. The mass of a C-12 atom is 12 a.m.u. a.m.u. stands for a tomic m ass u nit Formula Mass – The sum of the atomic masses of all the atoms in a compound. Section 10-1.

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Chapter 10

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  1. Chapter 10 The Mole

  2. Section 10-1 • Atomic Mass– The mass of an atom relative to the mass of C-12. • The mass of a C-12 atom is 12 a.m.u. • a.m.u. stands for atomic mass unit • Formula Mass – The sum of the atomic masses of all the atoms in a compound.

  3. Section 10-1 • Formula Mass for Water (H2O) • 1. Determine the number of atoms of each element. • H = 2 • O = 1 • 2. Multiply the # of atoms by the atomic mass of the element. • H: 2 x 1.00 amu = 2.00 amu • O: 1 x 15.99 amu = 15.99 amu • 3. Add results together • 2.00 amu + 15.99 amu = 17.99 amu – the formula mass for water.

  4. Section 10-1 • The Mole • Comes from the German word “molekularewight”, which means “molecular weight”. • In the early 1900’s scientists needed to find a way to relate the atomic mass of 1 atom of an element in atomic mass units to an amount of that element in grams. • A relationship between the mass of an atom and the quantity of the element was determined.

  5. Section 10-1 • 1 mole of any element or compound contains 6.02 x 1023 particles. • 6.02 x 1023 is called Avogadro’s number. • Amadeo Avogadro – (1776-1856) Italian chemist and physicist. • Element = 6.02 x 1023atoms. • Ionic compound = 6.02 x 1023formula units. • Covalent compound = 6.02 x 1023molecules.

  6. Section 10-1 • Molar Mass – the mass in grams of 1 mole of a substance. • Pg. 320, figure 10-14. • This depends on the masses of the particles, atoms, molecules, or formula units that make up the compound.

  7. Section 10-2 • Mole Conversions • The word “mole” represents an amount of something, just like the word “dozen” or “ton” or “gross” represent an amount of something. • A mole measures both mass and a number of particles, therefore, conversions can be made. g x (mol/g) = mole Mass moles x (g/mol) = g Mole

  8. Section 10-2 • Example: Convert 125.3 g CaCl2 to moles • 125.3g CaCl2 x 1 mol = 1.12 mol CaCl2 110.98g Now you try these! 1. convert 2.3 mol NaCl to grams 2. convert 35.8 g MgBr2 to moles

  9. Section 10-2 • A mole also tells you the number of particles of something you have. Mol x (6.02x1023/mole) = particles Particles Moles Particles x (mol/6.02x1023) = moles

  10. Section 10-2 • Example: A piece of marble contains 8.74x1023 formula units of Ca(CO3). How many moles is this? • 8.74x1023 x 1 mole = 1.45 moles Ca(CO3) 6.02x1023 Now you try Practice Problems 7 & 8 on page 327.

  11. Section 10-2 • Multi-step conversions mass particles mole

  12. Section 10-2 • As you know, not all chemicals are liquid or solid, some are gas, so we must also have a mole conversion for gases. • 1 mole of a gas occupies the same volume as 1 mole of any other gas at standard temperature and pressure (STP) • STP = 0° C and 1 atmosphere of pressure • At STP the molar volume of any gas is 22.4 L.

  13. Section 10-2 Mole-volume conversions Mole mole x (vol/mole) = volume volume x (mole/vol) = mole Volume at STP

  14. Section 10-2 • Everything is related! mass particles mole volume at STP

  15. Section 10-3 • Percent Composition – the mass of each element in a compound compared to the entire mass of the compound. • Handout #1 • Tells you the percent of the mass made up by each element in the compound. • Percentages of each must add up to 100%.

  16. Section 10-3 • Empirical formula – A formula that gives the simplest whole number ration of atoms in a compound. • Handout #1 • Molecular formula – the formula that gives the actual number of atoms of each element in a covalent compound • Handout #1 • It is the empirical formula x a whole number

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