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Chapter 11

Chapter 11. Chemical Quantities Chemistry Tracy Bonza Sequoyah High School. How do 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.

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Chapter 11

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  1. Chapter 11 Chemical Quantities Chemistry Tracy Bonza Sequoyah High School

  2. How do 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 Avogadro's number.

  4. 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.

  5. Representative particles • How many oxygen atoms in the following? • CaCO3 • Al2(SO4)3 • How many ions in the following? • CaCl2 • NaOH • Al2(SO4)3

  6. Molar 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.

  7. Potassium Zinc Neon 39g 65.4g 20.2g What is the molar mass of: What does that mean? It means that 39g of K is 1 mole of potassium How many particles of potassium is that? 6.02 x 10 23 See…….we can count particles without counting them ……….we weighed them!!!

  8. 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

  9. 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 Molar mass of CH4 is 16.05g • The mass of one mole of a molecular compound.

  10. Molar Mass • The generic term for the mass of one mole. • The same as gram molecular mass (for molecules/covalent compounds), gram formula mass (for formula units/ionic compounds), and gram atomic mass (for atoms). • I will use the generic molar mass (sometimes I slip up and call it gram formula mass so work with me here!!!!!)

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

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

  13. 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.

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

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

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

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

  18. 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

  19. 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

  20. 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

  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 • 1 mole NaOH = 40.00 g

  22. Examples • How much would 2.34 moles of carbon weigh, in grams? • How many moles of magnesium in 24.31 g of Mg? • How many moles is 4.56 g of CO2 ? • How many grams is 9.87 moles of H2O?

  23. Moles and Molecules • 6.02 x 1023 =1 mole too! • So we can convert from moles to molecules too! • Try it: • How many molecules are there in 3.6 moles of sodium hydroxide? • How many moles is 3.6x 1034 molecules of gold?

  24. 3 Step Problems • Converting from grams to molecules/atoms and molecules/atoms to grams • How many atoms of silver are in 2.3g of silver?

  25. Examples • How many molecules in 6.8 g of CH4? • 49 molecules of C6H12O6 weighs how much? • How many atoms of lithium in 1.00 g of Li? • How much would 3.45 x 1022 atoms of U weigh?

  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 • Avogadro’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 ?g Moles

  37. Mass Volume ?g Moles

  38. Mass Volume 22.4 L ?g Moles

  39. Mass Volume 22.4 L ?g Moles Representative Particles

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

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

  42. Mass Volume 22.4 L ?g Moles 6.02 x 1023 Representative Particles molecules 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 composition 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 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.

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