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Chapter 7

Chapter 7. Chemical Quantities or. "Our Friend the Mole". How do you measure how much?. You can measure mass, volume, or you can count pieces of a substance. We measure mass in grams. We measure volume in liters. We count pieces in MOLES. No, not that kind of mole!!!. Moles.

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Chapter 7

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  1. Chapter 7 Chemical Quantities or "Our Friend the Mole"

  2. How do you measure how much? • You can measure mass, volume, or you can count pieces of a substance. • We measure mass in grams. • We measure volume in liters. • We count pieces in MOLES. No, not that kind of mole!!!

  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. • The units of Avogadro’s number can be whatever particle you are counting. • Examples: atoms, molecules, ions, formula units, etc… • In chemistry these are called Representative Particles

  4. What are Representative Particles? • The smallest pieces of a substance. • For an element it is an atom. • Unless it is diatomic • For a covalent compound it is a molecule. • For an ionic compound it is a formula unit. • If it has a charge, it is an ion.

  5. How do We Use Moles? • Moles are used as conversion factors. • This means they are used to change units. • Remember, when solving using conversion factors there are 3 questions you want to ask yourself: • What unit do you want to get rid of? • Where does it go to cancel out? • What can you change it into?

  6. Calculation question 1 • How many molecules of CO2 are the in 4.56 moles of CO2 ? • Answer: 2.75 x 1024 molecules of CO2

  7. Calculation question • How many moles of water is 5.87 x 1022 molecules of water? • ANSWER: 0.0975 moles of water

  8. Measuring Moles • The atomic mass unit (amu) is 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: • The mass of the average atom in amu. • This mass is equivalent to the mass of 1 mole of those atoms in grams.

  9. 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.

  10. Examples • How much would 2.34 moles of carbon weigh? • 28.1 grams of C • How many moles of magnesium in 4.61 g of Mg? • 0.190 moles of Mg

  11. 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 their mass is by using the periodic table and then add them up.

  12. Example • 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

  13. Gram Molecular Mass or GMM • The mass of one mole of a molecular compound. Gram Formula Mass - The mass of one mole of an ionic compound. • Calculated the same way as GMM. • 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

  14. Molar Mass • The generic term for the mass of one mole. • The same as gram molecular mass, gram formula mass, and gram atomic mass. • This is the term we will be using in class. I do NOT weigh that much! The balance must be lying!!!

  15. Examples • Calculate the molar mass of the following and tell me what type it is (gmm, gfm, or gam). • Na2S • N2O4 • C • Ca(NO3)2 • C6H12O6 • (NH4)3PO4 78.05 g/mol 92.02 g/mol 12.01 g/mol 164.10 g/mol 180.18 g/mol 149.12 g/mol

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

  17. Molar Mass • The number of grams in 1 mole of atoms, formula units, or molecules. • We can make conversion factors from these. • It will allow us to change grams of a compound to moles of a compound. • Or moles to grams.

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

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

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

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

  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

  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 add these together to get the molar mass – use it to convert! • 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. 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

  26. Examples 0.104 mol CO2 • How many moles is 4.56 g of CO2 ? • How many grams is 9.87 moles of H2O? 177.86 g H2O

  27. Gases and the Mole Where is my air tank?? I hope it’s filled to 22.4 L

  28. Gases • Many of the chemicals we deal with are gases. • They are difficult to weigh, so we’ll measure volume. • We still need to know how many moles of gas we have. • Two things affect the volume of a gas: • Temperature and pressure. • So we have to compare at the SAME temp. and pressure.

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

  30. 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?

  31. We have learned how to: • change moles to grams • moles to atoms • moles to formula units • moles to molecules • moles to liters

  32. Mass Volume Representative Particles Molecules Atoms 22.4 L PT Moles 6.02 x 1023 Count

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

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

  35. Examples • Calculate the percent composition of C2H4. • Aluminum carbonate.

  36. Percent to Mass • Multiply % by the total mass to find the mass of that component. • How much aluminum in 450 g of aluminum carbonate?

  37. Empirical Formula From percentage to formula

  38. The Empirical Formula • The lowest whole number ratio of elements in a compound. • The molecular formula is 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

  39. Finding Empirical Formulas • Just find the lowest whole number ratio. • C6H12O6 , CH4N2 • It is not just the ratio of atoms, it is also the ratio of moles of atoms. Calculating Empirical Formulas • We can get ratio from percent composition. • Assume you have 100 g. • The percentages become grams. • Convert grams to moles. • Find lowest whole number ratio by dividing everything by the smallest moles.

  40. Example • Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N. • Assume 100 g so: • = 3.22 mole C • = 16.1 mole H • = 3.22 mole N

  41. Example Continued: • The smallest number of moles is 3.22 mol so divide your answers by 3.22 to get the mole ratio for the formula. • The ratio is: • So the formula is: • C1H5N1

  42. Additional Examples: • A compound is 73.9 % Hg and 26.1 % Cl. What is the empirical formula? • Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula? HgCl2 C4H5N2O

  43. Empirical to molecular • Since the empirical formula is the lowest ratio the actual molecule would weigh the same or more. • This difference in weight would be by a whole number multiple. • To find this multiple: divide the actual molar mass by the the mass of one mole of the empirical formula. • You will get a whole number. • Multiply the subscripts of the empirical formula by this number.

  44. Example • A compound has an empirical formula of ClCH2 and a molar mass of 98.96 g/mol. What is its molecular formula? • A compound has an empirical formula of CH2O and a molar mass of 180.0 g/mol. What is its molecular formula? Cl2C2H4 C6H12O6

  45. Percent Composition to Molecular Formula: • Ibuprofen is 75.69 % C, 8.80 % H, 15.51 % O, and has a molar mass of about 207 g/mol. What is its molecular formula? C13H18O2

  46. Example: • Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula? • C  0.4948*194g = 95.99 g C • = 8 mol C • H  0.0515*194g = 9.99 g H • = 10 mol H • N  0.2887*194g = 56.01 g N • = 4 mol N • O  0.1649*194g = 31.99 g O • = 2 mol O C8H10N4O2

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