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Chapter 6

Chapter 6. Chemical Quantities. 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. Moles. Defined as the number of carbon atoms in exactly 12 grams of carbon-12.

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Chapter 6

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  1. Chapter 6 Chemical Quantities

  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. 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.

  4. 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.

  5. 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

  6. 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?

  7. Measuring Moles • Remember relative atomic mass? • The amu was one twelfth the mass of a carbon 12 atom. • Since the mole is the number of atoms in 12 grams of carbon-12, • the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.

  8. Gram Atomic Mass • The mass of 1 mole of an element in grams. • 12.01 grams of carbon has the same number of pieces as 1.008 grams of hydrogen and 55.85 grams of iron. • We can write this as 12.01 g C = 1 mole • We can count things by weighing them.

  9. 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?

  10. 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

  11. 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.

  12. Gram Formula Mass • The mass of one mole of an ionic compound. • Calculated the same way. • What is the GFM of Fe2O3? • 2 moles of Fe x 55.85 g = 111.70 g • 3 moles of O x 16.00 g = 48.00 g • The GFM = 111.70 g + 48.00 g = 159.70g

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

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

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

  16. 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.

  17. For example • How many moles is 5.69 g of NaOH?

  18. For example • How many moles is 5.69 g of NaOH?

  19. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles

  20. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH

  21. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g

  22. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  23. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  24. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  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 then 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. Percent Composition • Like all percents • Part x 100 % whole • Find the mass of each component, • divide by the total mass.

  44. Example • Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

  45. Getting it from the formula • If we know the formula, assume you have 1 mole. • Then you know the pieces and the whole.

  46. Examples • Calculate the percent composittion of C2H4? • Aluminum carbonate.

  47. Empirical Formula From percentage to formula

  48. The Empirical Formula • The lowest whole number ratio of elements in a compound. • The molecular formula = the actual ratio 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. Calculating Empirical • Means we can get ratio from percent composition. • Assume you have a 100 g. • The percentages become grams. • Can turn grams to moles. • Find lowest whole number ratio by dividing by the smallest.

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