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Chapter 8

Chapter 8. Chemical Composition. How you measure how much?. You can measure mass, or volume, or you can count pieces. We measure mass in grams. We measure volume in liters. We count pieces in MOLES. VERY. A large amount!!!!. A. What is the Mole?.

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Chapter 8

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  1. Chapter 8 Chemical Composition

  2. How you measure how much? • You can measure mass, • or volume, • or you can count pieces. • We measure mass in grams. • We measure volume in liters. • We count pieces in MOLES.

  3. VERY A large amount!!!! A. What is the Mole? • A counting number (like a dozen) • Avogadro’s number (NA) • 1 mol = 6.02  1023 items

  4. What is the Mole? • 1 mole of pennies would cover the Earth 1/4 mile deep! HOW LARGE IS IT??? • 1 mole of hockey pucks would equal the mass of the moon! • 1 mole of basketballs would fill a bag the size of the earth!

  5. Moles • Defined as the number of carbon atoms in exactly 12 grams of carbon-12. • 1 mole is 6.02 x 1023 particles. • Treat it like a very large dozen • 6.02 x 1023 is called Avagadro’s number.

  6. Representative particles • The smallest pieces of a substance. • For a molecular compound it is a molecule. • For an ionic compound it is a formula unit. • For an element it is an atom.

  7. Types of questions • How many oxygen atoms in the following? • CaCO3 • Al2(SO4)3 • How many ions in the following? • CaCl2 • NaOH • Al2(SO4)3

  8. Molar Mass • Mass of 1 mole of an element or compound. • Atomic mass tells the... • atomic mass units per atom (amu) • grams per mole (g/mol) • Round to 2 decimal places

  9. Molar Mass Examples • carbon • aluminum • zinc 12.01 g/mol 26.98 g/mol 65.39 g/mol

  10. Molar Mass Examples • water • sodium chloride • H2O • 2(1.01) + 16.00 = 18.02 g/mol • NaCl • 22.99 + 35.45 = 58.44 g/mol

  11. Molar Mass Examples • sodium bicarbonate • sucrose • NaHCO3 • 22.99 + 1.01 + 12.01 + 3(16.00) = 84.01 g/mol • C12H22O11 • 12(12.01) + 22(1.01) + 11(16.00) = 342.34 g/mol

  12. MASS IN GRAMS MOLES NUMBER OF PARTICLES Molar Conversions molar mass 6.02  1023 (g/mol) (particles/mol)

  13. Molar Conversion Examples • How many moles of carbon are in 26 g of carbon? 26 g C 1 mol C 12.01 g C = 2.2 mol C

  14. Molar Conversion Examples • How many molecules are in 2.50 moles of C12H22O11? 6.02  1023 molecules 1 mol 2.50 mol = 1.51  1024 molecules C12H22O11

  15. Molar Conversion Examples • Find the mass of 2.1  1024 molecules of NaHCO3. 2.1  1024 molecules 1 mol 6.02  1023 molecules 84.01 g 1 mol = 290 g NaHCO3

  16. Types of questions • How many molecules of CO2 are the in 4.56 moles of CO2 ? • How many moles of water is 5.87 x 1022 molecules? • How many atoms of carbon are there in 1.23 moles of C6H12O6 ? • How many moles is 7.78 x 1024 formula units of MgCl2?

  17. Examples • How much would 2.34 moles of carbon weigh? • How many moles of magnesium in 24.31 g of Mg? • How many atoms of lithium in 1.00 g of Li? • How much would 3.45 x 1022 atoms of U weigh?

  18. What about compounds? • in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atoms • To find the mass of one mole of a compound • determine the moles of the elements they have • Find out how much they would weigh • add them up

  19. What about compounds? • What is the mass of one mole of CH4? • 1 mole of C = 12.01 g • 4 mole of H x 1.01 g = 4.04g • 1 mole CH4 = 12.01 + 4.04 = 16.05g • The Gram Molecular mass of CH4 is 16.05g • The mass of one mole of a molecular compound.

  20. Gram Formula Mass • The mass of one mole of an ionic compound. • Calculated the same way.

  21. Molar Mass • The generic term for the mass of one mole. • The same as gram molecular mass, gram formula mass, and gram atomic mass.

  22. Examples • Calculate the molar mass of the following and tell me what type it is. • Na2S • N2O4 • C • Ca(NO3)2 • C6H12O6 • (NH4)3PO4

  23. Using Molar Mass Finding moles of compounds Counting pieces by weighing

  24. Molar Mass • The number of grams of 1 mole of atoms, ions, or molecules. • We can make conversion factors from these. • To change grams of a compound to moles of a compound.

  25. Examples • How many moles is 4.56 g of CO2 ? • How many grams is 9.87 moles of H2O? • How many molecules in 6.8 g of CH4? • 49 molecules of C6H12O6 weighs how much?

  26. Gases and the Mole

  27. Gases • Many of the chemicals we deal with are gases. • They are difficult to weigh. • Need to know how many moles of gas we have. • Two things effect the volume of a gas • Temperature and pressure • Compare at the same temp. and pressure.

  28. Standard Temperature and Pressure • 0ºC and 1 atm pressure • abbreviated STP • At STP 1 mole of gas occupies 22.4 L • Called the molar volume • Avagadro’s Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles.

  29. Examples • What is the volume of 4.59 mole of CO2 gas at STP? • How many moles is 5.67 L of O2 at STP? • What is the volume of 8.8g of CH4 gas at STP?

  30. Density of a gas • D = m /V • for a gas the units will be g / L • We can determine the density of any gas at STP if we know its formula. • To find the density we need the mass and the volume. • If you assume you have 1 mole than the mass is the molar mass (PT) • At STP the volume is 22.4 L.

  31. Examples • Find the density of CO2at STP. • Find the density of CH4 at STP.

  32. The other way • Given the density, we can find the molar mass of the gas. • Again, pretend you have a mole at STP, so V = 22.4 L. • m = D x V • m is the mass of 1 mole, since you have 22.4 L of the stuff. • What is the molar mass of a gas with a density of 1.964 g/L? • 2.86 g/L?

  33. All the things we can change

  34. We have learned how to • change moles to grams • moles to atoms • moles to formula units • moles to molecules • moles to liters • molecules to atoms • formula units to atoms • formula units to ions

  35. Mass Moles

  36. Mass PT Moles

  37. Mass Volume PT Moles

  38. Mass Volume 22.4 L PT Moles

  39. Mass Volume 22.4 L PT Moles Representative Particles

  40. Mass Volume 22.4 L PT Moles 6.02 x 1023 Representative Particles

  41. Mass Volume 22.4 L PT Moles 6.02 x 1023 Representative Particles Atoms

  42. Mass Volume 22.4 L PT Moles 6.02 x 1023 Representative Particles Ions Atoms

  43. Percentage Composition • the percentage by mass of each element in a compound

  44. 127.10 g Cu 159.17 g Cu2S 32.07 g S 159.17 g Cu2S Percentage Composition • Find the % composition of Cu2S.  100 = %Cu = 79.852% Cu %S =  100 = 20.15% S

  45. 28 g 36 g 8.0 g 36 g Percentage Composition • Find the percentage composition of a sample that is 28 g Fe and 8.0 g O.  100 = 78% Fe %Fe =  100 = 22% O %O =

  46. Percentage Composition • How many grams of copper are in a 38.0-gram sample of Cu2S? Cu2S is 79.852% Cu (38.0 g Cu2S)(0.79852) = 30.3 g Cu

  47. 36.04 g 147.02 g Percentage Composition • Find the mass percentage of water in calcium chloride dihydrate, CaCl2•2H2O? 24.51% H2O %H2O =  100 =

  48. The Empirical Formula • The lowest whole number ratio of elements in a compound. • The molecular formula the actual ration of elements in a compound. • The two can be the same. • CH2 empirical formula • C2H4 molecular formula • C3H6 molecular formula • H2O both

  49. Calculating Empirical • Just find the lowest whole number ratio • C6H12O6 • CH4N • It is not just the ratio of atoms, it is also the ratio of moles of atoms. • In 1 mole of CO2there is 1 mole of carbon and 2 moles of oxygen. • In one molecule of CO2 there is 1 atom of C and 2 atoms of O.

  50. CH3 Empirical Formula • Smallest whole number ratio of atoms in a compound C2H6 reduce subscripts

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