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Chapter 7

Chapter 7. The Mole and Chemical Compostion. Unit Essential Question:. How can chemical composition be determined?. Section 1 Lesson Essential Question:. Note: the majority of this section is review!. How is the mole used in conversions?. Section 1: Avogadro ’ s Number and Molar Conversions.

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Chapter 7

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  1. Chapter 7 The Mole and Chemical Compostion

  2. Unit Essential Question: How can chemical composition be determined?

  3. Section 1 Lesson Essential Question: Note: the majority of this section is review! How is the mole used in conversions?

  4. Section 1: Avogadro’s Number and Molar Conversions • 1 mole = 6.022 x 1023 particles • SI unit for amount of substance. • It’s a counting unit like a dozen, ream, case, etc. • Remember that the unit of particles can be: ions, molecules (mcs.), atoms, formula units (f.u.), etc. covalent compounds ionic compounds

  5. Recall your mole map!

  6. Converting moles  particles • Same as Chapter 3, but it will involve molecules, formula units, or ions instead of just atoms. • Steps: 1) Need 1mol = 6.022 x1023 molecules, etc. 2) Use dimensional analysis- turn this into a fraction! *Be sure to place the correct units on the top and bottom so they cancel!

  7. Sample Problems 1 & 2: Moles & Particles • Find the number of molecules in 2.5 mol of sulfur dioxide. • A sample contains 3.01 x 1023 molecules of sulfur dioxide. Determine the amount in moles.

  8. Molar Mass • Amount of mass (in grams) in 1 mole of a substance. • Use molar masses from the periodic table. • Round to 2 decimal places! • Use units of g/mol. • Example: • C: 12.01g/mol means that 1 mol C = 12.01 g • Cl: 35.45g/mol means that 1 mol Cl = 35.45g • Use to convert between moles and mass.

  9. Sample Problems 3 & 4: Moles & Mass • What is the mass of 5.3mol Be? • If you have 27.0g of manganese, how many moles do you have?

  10. Molar Masses of Compounds • Add together the molar masses of all elements or ions present. • Ex: CH4 • C: 12.01g/mol H: 1.01g/mol • 12.01g/mol + 4(1.01g/mol) = 16.05g/mol • This means that 1 mole of CH4 has a mass of 16.05g. • You will need to calculate the molar mass of a compound whenever you are converting between mass and moles!

  11. Additional Molar Mass Examples: • Element • Ag = 107.87 g/mol • Diatomic Element/molecule • Br2 = 79.90 x 2 = 159.80 g/mol • Molecule (Covalent compound) • H2O = (1.01 x 2) + 16.00 = 18.02 g/mol • Formula unit (Ionic compound) • Ca(NO3)2 = 40.08 + (2 x 14.01) + (6 x 16.00) = 164.10 g/mol

  12. Sample Problem 5: Mass to Moles with a Compound • Find the number of moles present in 47.5 g of glycerol, C3H8O3. • Hint: you will need to calculate the molar mass of glycerol!

  13. Sample Problem 6: Number of Particles to Mass • Remember- you can’t go directly between mass (g) and the number of particles! You must convert to moles first! • Find the mass in grams of 2.44 x 1024 atoms of carbon.

  14. Unit Essential Question: How can chemical composition be determined?

  15. Lesson Essential Question: How are average atomic masses calculated?

  16. Section 2: Relative Atomic Mass and Chemical Formulas • Periodic table masses are averages of all isotopes present. • Recall that we said a weighted average is used- takes into account the amount of each isotope. • Average atomic mass: (% x atomic mass)+(% x atomic mass)+… 100 • Note: % is the percent abundance (how often the element is found as that isotope in nature).

  17. Sample Problem • The mass of a Cu-63 atom is 62.94 amu, and that of a Cu-65 atom is 64.93 amu. If the abundance of Cu-63 is 69.17% and the abundance of Cu-65 is 30.83%, what is the average atomic mass of copper?

  18. Lesson Essential Questions: What information can be determined from formulas?How can formulas be determined?

  19. Calculating Percent Composition • Tells you the percent each element makes up of the whole compound. Step 1: Determine the molar mass of the entire compound. Step 2: Divide each element’s total molar mass by the molar mass of the compound. Step 3: Multiply by 100 to get percent. Step 4: Check your answer by adding up the percentages to makes sure they equal 100%.

  20. Section 3: Formulas and Percent Composition • Can verify percent composition of a compound to determine formula/identity. • Example: • Iron and oxygen form two compounds: • Fe2O3 and FeO • Fe2O3 = 69.9% Fe and 30.1% O • FeO = 77.7% Fe and 22.3% O

  21. Sample Problem #I • Calculate the percent composition of copper (I) sulfide. • Calculate the percent composition of isopropyl alcohol, (CH3)2CHOH. Sample Problem #2

  22. Determining Empirical Formulas • The empirical formula shows the simplest ratio of elements in the compound. • Given percent composition data, you can determine the empirical formula of a compound. Step 1: Assume 100 g of the sample- turn %’s into grams. Ex: 18.2% O  18.2g Step 2: Convert grams to moles. Step 3: Divide each mole value by the smallest mole value. This will tell you the number of each element that appears in the formula.

  23. Determining Empirical Formulas Cont. Step 4: If you get a decimal, multiply ALL numbers by a whole number to turn the decimal into a whole number. • The numbers you will need to multiply by should be relatively small (2, 3, etc.)

  24. Sample Problem #1 • Chemical analysis of a liquid shows that it is 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula of this substance. • A compound is found to contain 38.77% Cl and 61.23% O. What is the empirical formula? Sample Problem #2

  25. Molecular Formulas • Show the actual numbers of elements in the compound- not necessarily the simplest formula. • They will be a whole number multiple of the empirical formula (can’t be a decimal). • In other words: n(empirical formula) = molecular formula where n is a whole number. • Ex: 6(CH2O)  C6H12O6 • Molecular and empirical formulas can be the same!

  26. Molecular Formulas Cont.

  27. Molecular Formulas Cont. • The molecular formula can be determined from the empirical formula and experimental molar mass of a compound. Step 1: Determine the molar mass of the given empirical formula. Step 2: Solve for n by dividing the experimental molar mass by the molar mass of the empirical formula. *Remember: n(empirical formula) = molecular formula Step 3: Multiply the subscripts in the empirical formula by n.

  28. Sample Problem #1 • The empirical formula for a compound is P2O5. Its experimental molar mass is 284g/mol. Determine the molecular formula of the compound. • A brown gas has the empirical formula NO2. Its experimental molar mass is 46g/mol. What is the molecular formula? Sample Problem #2

  29. Hydrates • Not in the textbook. • Hydrates – ionic compounds that contain water molecules within the crystal structure. • Example: CuSO4•5H2O • Anhydrous – without the water = CuSO4

  30. Hydrates Cont. • Can be calculated if given the masses of the hydrate and anhydrous mass and the formula of the ionic compound. Step 1: Determine the mass of water in the hydrate (subtract anhydrous mass). Step 2: Convert the anhydrous ionic compound mass and water mass to moles. Step: Divide both molar amounts by the smallest number. This gives you the number of water molecules in the hydrate.

  31. Sample Problem #1 • A 5.82 g sample of Mg(NO3)2·XH2O in an evaporating dish is heated until it is dry.  The mass of the anhydrous sample is 2.63 g Mg(NO3)2.  What is the formula for the hydrate? • What percentage, by mass, of water is found in the hydrate CuSO4·5H2O? Sample Problem #2

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