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Chemical Composition Chapter 8

Chemical Composition Chapter 8. The Mole. A mole is a counting unit, just like a dozen, except that a mole is much larger. A mole (mol) is a number that represents the number of representative particles of a substance.

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Chemical Composition Chapter 8

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  1. Chemical CompositionChapter 8

  2. The Mole • A mole is a counting unit, just like a dozen, except that a mole is much larger. • A mole (mol) is a number that represents the number of representative particles of a substance. • It is always equal to 6.02 x 1023 and is referred to as Avogadro’s number.

  3. All these samples of pure elements contain the same number (a mole) of atoms: 6.022 x 1023 atoms.

  4. One-mole samples of iron (nails), iodine crystals, liquid mercury, and powdered sulfur.

  5. A 1-mol sample of graphite (a form of carbon) weighs 12.01 g.

  6. Representative Particles • Atoms - single element, for example: Na, C • Molecule (includes diatomic molecules) – contains no metals, for example: H2O, C12H22O11, HCl, N2, O2, F2, CO2 • Ion – only ions, for example, Ca2+,SO4-2, P3- • Formula Unit – always starts with a metal, for example: CaF2, CuSO4

  7. Using Avogadro’s Number • Avogadro’s # can be used as a conversion factor where 1 mol= 6.02 x 1023 atoms, ions, molecules, or formula units depending on the type of substance you have.

  8. Using Avogadro’s Number • Convert 4.27 x 1023 molecules of CO2 to moles • How many formula units are there in 2.78 moles of CaCl2?

  9. Using Avogadro’s Number • How many molecules are there in 0.85 moles of octane, C8H18? • How many moles of Na+ are there in 7.6 x 1022 Na+ ions?

  10. Various numbers of methane molecules showing their constituent atoms.

  11. Various numbers of methane molecules showing their constituent atoms

  12. More Mole Conversions • How many formula units of Al(NO3)3 are there in 0.750 moles of Al(NO3)3? How many atoms of oxygen are in the sample of Al(NO3)3?

  13. Molar Mass Conversions • When converting between grams and moles, 1 mole of an element equals the average atomic mass of that element in grams, rounded to one number after the decimal. • 1 mol Ca = 40.1 g Ca 1 mol P = 31.0 g P • 1 mol Cl = 35.5 g Cl 1 mol K = 39.1 g K • 1 mol S = 32.1 g S 1 mol Cu = 63.6 g Cu • 1 mol Br = 79.9 g Br 1 mol Si = 28.1 g Si

  14. Molar Mass Conversions • How many moles are there in a 10.0 gram sample of Al? • How many grams are there in a 0.643 mol sample of Fe?

  15. Grams and Avogadro’s # Conversions • How many grams of beryllium are present in a 4.18 x 1024 atom sample of Be? • How many atoms of chromium are present in a 23.8 gram sample of Cr?

  16. Determining Molar Mass • Molar mass – mass of 1 mol of a pure substance usually written in grams per mole (g/mol) • The subscript next to each element in a formula tells you how many moles of that particular element are in the compound.

  17. Determining Molar Mass • When finding molar mass of a compound if there are parentheses in the formula, you must multiply each element in the parentheses times the number outside to determine the total number of moles of each element.

  18. Molar Mass Determination • Determine the molar mass for: • Fe2(SO4)3 • Cu(ClO3)2 • (NH4)2C4H4O6 .3H2O

  19. Molar Mass & Avogadro’s # Conversions with Compounds • Determine the number of moles in a 22.3 gram sample of Al2(CO3)3

  20. Molar Mass & Avogadro’s # Conversions with Compounds • How many grams of ammonium phosphate, (NH4)3PO4, are present in a 0.082 mole sample of (NH4)3PO4?

  21. Molar Mass & Avogadro’s # Conversions with Compounds • How many grams of zinc chloride, ZnCl2, are present in a 8.21 x 1022 formula unit sample of ZnCl2?

  22. Molar Mass & Avogadro’s # Conversions with Compounds • Determine the number of formula units are there in a 47.9 gram sample of Mg(NO3)2 • How many atoms of nitrogen are in this sample of Mg(NO3)2?

  23. Molar Volume of a Gas • The volume of a gas varies with changes in temperature and pressure. Usually though, the volume of a gas is measured at a standard temperature and pressure. • Standard Temperature and Pressure (STP) – 1 atmosphere of pressure (1 atm) and 0oC (273 K). • At STP, 1 mole of any gas occupies 22.4 L. • Molar volume of gas– @ STP 1 mol gas= 22.4 L.

  24. Molar Volume Problems • At STP, how many moles of oxygen gas, O2, are present in a 47.3 L sample of O2? • How many liters of nitrogen gas, N2, are in a 2.75 mole sample of N2 at STP?

  25. Molar Volume Problems • At STP, how many liters of nitrogen gas, N2, are present in 4.81 x 1023 molecules of N2 gas? • How many grams of CO2 gas are present in a 2.00 L bottle of coca-cola at STP?

  26. Percent Composition of a Compound • To determine percent composition • Determine molar mass of your compound (g/mol). • Divide the total mass for each element by the molar mass of the compound. • Multiply each element’s % by 100. If you add up the percentages for all of the elements, you should get very close to 100%.

  27. % Composition Example • Determine the percent composition for each element in the following compounds: • K2SO4 • Ba3(PO4)2. 2H2O

  28. Empirical Formula vs. Molecular Formulas • Empirical formula – lowest whole-number ratio of the moles in a compound. • Molecular formulas are simple multiples of empirical formulas that can be reduced if each element in the compound is divisible by the same whole number. • Empirical formulas cannot be reduced any further and still have whole numbers for each element in the compound.

  29. Empirical Formulas vs. Molecular Formulas Examples • C6H9O3 • C12H22O11 • C4H6O2 • C3H8N4 • Ca3(S2O3)2

  30. Determining Empirical Formulas • If given percentages for each element, convert each one directly to grams. • Divide the grams of each element by its molar mass to get moles of each element. • Divide each number of moles by the smallest number of moles.

  31. Determining Empirical Formulas • All answers must be within 0.1 of a whole number. If not, then all of the elements must be multiplied by a small whole number so that all elements are whole numbers. If 0.5, then multiply all elements by 2 If 0.33 or 0.67, multiply all elements by 3 If 0.25 or 0.75, multiply all elements by 4

  32. Empirical Formula Examples • A compound is composed of 50.0% S and 50.0% O. What is the empirical formula of this compound?

  33. Empirical Formula Determination Examples • A compound contains 33.88% Cu, 14.94% N, and the remaining % is Oxygen. What is the empirical formula of this compound?

  34. Determine the Empirical Formula of a compound that contains 55.2% Carbon, 2.3% Hydrogen , 18.4% Oxygen and 24.1% Nitrogen

  35. Determining Molecular Formulas • Determine the empirical formula for the compound. • Determine the molar mass of the empirical formula. • Divide the molar mass of the molecular formula over the molar mass of the empirical formula. • Multiply the answer from #3 by the empirical formula.

  36. Determining Molecular Formulas Molar Mass of Molecular Formula ---------------------------- = X Molar Mass of Empirical Formula Molecular Formula = (Empirical formula)X • To check to see if your molecular formula is correct, determine the molar mass of your molecular formula. It should be equal to the molar mass of the molecular formula you were given in the problem.

  37. Molecular Formula Examples • Example: A compound has an empirical formula of C2H3O. What is the molecular formula if the molecular formula mass is 172.0 g/mol? • Determine the empirical and molecular formulas for a compound containing 94.1%O and 5.9% H. The molar mass of the molecular formula equals 34.0 g/mol.

  38. Molecular Formula Examples • Determine the empirical and molecular formulas for a compound containing 26.1%C, 50.7% N and 23.2% O. The molar mass of the molecular formula equals 828.0 g/mol.

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