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Ch. 6 Notes -- Chemical Composition What is a mole?

Ch. 6 Notes -- Chemical Composition What is a mole? Mole is a unit of quantity. Like a _____________. 1 mole = 6.02x10 23 atoms or molecules = “X” grams = 22.4 L gas. Dozen. The Mole!!! A counting unit Similar to a dozen, except instead of 12, it’s 602 billion trillion…

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Ch. 6 Notes -- Chemical Composition What is a mole?

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  1. Ch. 6 Notes -- Chemical Composition What is a mole? Mole is a unit of quantity. Like a _____________. 1 mole = 6.02x1023 atoms or molecules = “X” grams = 22.4 L gas Dozen

  2. The Mole!!! • A counting unit • Similar to a dozen, except instead of 12, it’s 602 billion trillion… (602,000,000,000,000,000,000,000) • ___________ (in scientific notation) • This number is named in honor of Amedeo_________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present…6.02 x 1023 6.02 x 10 23 Avogadro

  3. Just How Big is a Mole? • Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles. • If you had Avogadro's number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. • If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.

  4. The Mole 12 6.02 X 1023 • 1 dozen cookies = ___ cookies • 1 mole of cookies = ___________ cookies • 1 dozen cars = ___ cars • 1 mole of cars = __________ cars • 1 dozen Al atoms = ___ Al atoms • 1 mole of Al atoms = __________ atoms Note that the NUMBER is always the same, but the ______ is very different! Mole is abbreviated ______ . 12 6.02 X 1023 12 6.02 X 1023 MASS mol

  5. The Mole and Mass the sum • Mass in grams of 1 mole equal to __________ of the atomic masses Practice problem: Calculate the mass of 1 mole of CaCl2 Ca = 1 x ________ g/mol = 40.1 g/mol Cl = 2 x ________ g/mol = 71.0 g/mol 40.1 g/mol + 71.0 g/mol = __________ g/mol CaCl2 1 mole of CaCl2 = 111.1 g/mol 40.1 35.5 111.1

  6. 2 8 9 4 Ch. 7 Notes -- Chemical Quantities Practice Problems: (1) How many atoms of hydrogen are there in each compound? a) Ca(OH)2 ___ b) C3H8O___ c) (NH4)2HPO4 ___ d) HC2H3O2 ___ (2) Calculate the formula mass of each compound. (Add up all the atomic masses for each atom from the Periodic Table.) a) CaCO3 b) (NH4)2SO4 c) C3H6O d) Br2 2 N’s = 2 x 14.0 = 28.0 8 H’s = 8 x 1.0 = 8.0 S = 32.1 4 O’s = 4 x 16.0 = 64.0 Ca = 40.1 C = 12.0 3 O’s =3 x 16.0 = 48.0 Add them up! 132.1 g/mol Add them up! 100.1 g/mol 159.8 g/mol C = 3 x 12.0 = 36.0 H = 6 x 1.0 = 6.0 O =16.0 2 Br’s = 2 x 79.9 = Add them up! 58.0 g/mol

  7. 1 mole SO3 835 g SO3 10.4 moles of SO3 x = 80.1 g SO3 3) Convert 835 grams of SO3 to moles. 4) How many molecules of CH4 are there in 18 moles? 5) How many grams of helium are there in 5.6 x 1023 atoms of helium? 6) How many molecules are there in 3.7 grams of H2O? 6.02 x 1023 molecules CH4 18 moles CH4 1.08 x 1025 molecules CH4 x = 1 mole CH4 4.0 grams He 5.6 x 1023 atoms He x 3.72 grams He = 6.02 x 1023 atoms He 6.02 x 1023 molecules H2O 3.7 grams H2O 1.23 x 1023 molecules H2O x = 18.0 grams H2O

  8. Calculating Percent Composition by Mass Step 1: Find the formula mass of the compound by adding the individual masses of the elements together. Step 2: Divide each of the individual masses of the elements by the formula mass of the compound. Step 3: Convert the decimal to a % by multiplying by 100. Practice Problems: (1) Find the % composition of the elements in each compound. a) Na3PO4 b) SnCl4 = 0.421 = 42.1% 3 Na’s = 3 x 23.0 = 69.0 P = 31.0 4 O’s = 4 x 16.0 = 64.0 ÷ 164 Sn = 118.7 4 Cl’s = 4 x 35.5 = 142.0 ÷ 260.7 = 45.5% = 0.189 = 18.9% + ÷ 164 ÷ 260.7 = 54.5% 260.7 = 0.390 = 39.0% + ÷ 164 164

  9. Elements in the Universe: % Composition by Mass

  10. Earth’s Crust: % Composition by Mass

  11. Entire Earth (Including Atmosphere): % Composition by Mass

  12. Human Body: % Composition by Mass

  13. Meet The Elements

  14. Determining the Empirical Formula for a Compound whole ratio • The empirical formula for a compound is the simplest __________ number __________ of the atoms in the compound. Examples:H2O is the empirical formula for water. _______ is the empirical formula for glucose, C6H12O6. Practice Problems: What is the empirical formula for the following compounds? a) C6H6= ________ b) C8H14O2 = ________ c) C10H14O2 = _________ d) Ca5Br10 = ________ e) N3O9 = ________ C1H2O1 CH C4H7O C5H7O CaBr2 NO3

  15. Determining the Molecular Formula for a Compound whole # multiple • The molecular formula for a compound is either the same as the empirical formula ratio or it is a “_________ _________ of this ratio. It represents the true # of atoms in the molecule. Examples:1) H2O is the empirical & molecular formula for water. 2) CH2O is the empirical formula for sugar, ethanoic acid, and methanol. The molecular formula for glucose is C6H12O6, (___times the empirical ratio!) Practice Problems: (1) If the empirical formula for a compound is CH2, which of the following is a possible molecular formula for the compound? a) C8H16 b) C8H8 c) C4H2 d) C3H9 (2) If the empirical formula for a compound is C2H3, which of the following is a possible molecular formula for the compound? a) C2H6 b) C10H15 c) C6H12 d) C8H14 6

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