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Chemistry Chapter 7 – Chemical Quantities

Chemistry Chapter 7 – Chemical Quantities. Section 7.1 The Mole: A Measurement of Matter. Objectives: Describe how Avogadro’s number is related to a mole of any substance Calculate the mass of a mole of any substance. Representative Particles. Types of representative particles:

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Chemistry Chapter 7 – Chemical Quantities

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  1. ChemistryChapter 7 – Chemical Quantities

  2. Section 7.1 The Mole: A Measurement of Matter • Objectives: • Describe how Avogadro’s number is related to a mole of any substance • Calculate the mass of a mole of any substance

  3. Representative Particles • Types of representative particles: • Atoms (Single element) • Molecules (Covalent/Molecular Compound) • Formula Units (Ionic Compound) • Ions (+/- charge)

  4. Ex. Label each representative particle. • KF formula unit • CO2 molecule • He atom • SO42- ion • Ba(OH)2 formula unit • F2 molecule • H2O molecule • H3O+ ion

  5. What Is a Mole? • Mole(mol) – a quantity which represents 6.02 x 1023 representative particles of any given substance. • Avogadro’s Number – 6.02 x 1023 or 1 mole • The term “mol” is similar to: dozen, ream, bushel • Representative particle – the species present in a substance: atoms, molecules, formula units, ions.

  6. How large is a mol? • A mol of golf balls: • lined up would go to the sun and back ~1 billion times (dist to sun is ~92,000,000 miles) • A mol of animal moles: • spread over the Earth would make a layer 8 million animal moles thick

  7. Conversion Factors • 1 mole = 6.02 x 1023 atoms • 1 mole = 6.02 x 1023 molecules • 1 mole = 6.02 x 1023 formula units • 1 mole = 6.02 x 1023 ions

  8. Ex: How many atoms are there in 1.14 mol Ag? 1.14 mole Ag x 6.02 x 1023 atoms = 6.8628 x 1023 atoms Ag 1 mole Ag 6.86 x 1023 atoms Ag

  9. Ex: How many moles of magnesium is 1.25 x 1023 atoms of magnesium? 1.25 x 1023 atom Mg x 1 mole Mg = 0.207641196 mol Mg 6.02 x 1023 atoms Mg 0.208 moles Mg

  10. Ex: How many molecules is 0.360 mol of water? 0.360 mole H2O x 6.02 x 1023 molecules = 2.1672 x 1023 molecules H2O 1 mole H2O 2.17 x 1023 molecules H2O

  11. Ex: How many moles of NO2 are there in 4.65 x 1024 molecules of NO2? 4.65 x 1024 molecules NO2 x 1 mole NO2= 7.724 mol NO2 6.02 x 1023 molecules NO2 7.72 moles NO2

  12. Caution! • Be careful when being asked to convert moles of a compound into atoms! • We will need to multiply the final answer by the number of atoms in the compound. MOLE CROSSING

  13. Ex: How many atoms are there in 1.14 mol SO3? 1.14 mol x 6.02 x 1023 molecules x 4 atoms = 2.74512 x 1024 atoms 1 mole 1 molecule 2.75 x 1024 atoms

  14. Ex: How many atoms are in 2.12 mol of propane (C3H8)? 2.12 mol x 6.02 x 1023 molecules x 11 atoms = 1.403864 x 1025 atoms 1 mole 1 molecule 1.40 x 1025 atoms

  15. The Mass of a Mole of an Element • The gram atomic mass (gam) is the atomic mass of an element expressed in grams. We will use the periodic table to determine this. • Gram atomic mass of Carbon = 12.01 g • Gram atomic mass of Nitrogen = 14.01g • Gram atomic mass of Sulfur = 32.06 g • The gram atomic mass is equivalent to one mole of the atom. Use the Periodic Table

  16. The Mass of a Mole of a Compound • The gram molecular mass (gmm) of any molecular compound is the mass of 1 mole of that compound. We will again use the periodic table to determine this. • Find the gram molecular mass of the following: • H2O2 • N2O5 • Ca(OH)2

  17. The Mass of a Mole of a Compound • The mass of one mole of an ionic compound is the gram formula mass (gfm). A gram formula mass is calculated the same way as a gram molecular mass. • Find the gram formula mass of the following: • CaI2 • (NH4)2CO3

  18. Section 7.1 The Mole: A Measurement of Matter • Did We Meet Our Objectives? • Describe how Avogadro’s number is related to a mole of any substance • Calculate the mass of a mole of any substance Charlotte would be proud!

  19. Section 7.2 Mole-Mass and Mole-Volume Relationships • Objectives: • Use the molar mass to convert between mass and moles of a substance • Use the mole to convert among measurements of mass, volume, and number of particles

  20. The Mass of a Mole of an Element • Molar mass – mass of 1 mol of any substance.Can be used in calculations involving elements, molecular compounds, and ionic compounds • 1.0 mol of C has a mass of 12.01 g 12.01 g/mol • 1.0 mol of H2 has a mass of 2.02 g 2.02 g/mol • 1.0 mol H2O has a mass of 18.02 g 18.02 g/mol

  21. Ex: Find the mass, in grams, of 2.5 molsof Na. 2.5 mol Na x22.99 g Na= 57.475 g Na 1 mole Na 57 g Na

  22. Ex: Find the number of mols in 75.0 g of dinitrogen trioxide (N2O3). 75.0 g N2O3x1 mol N2O3= 0.98658 mol N2O3 76.02 g N2O3 0.987 mol N2O3

  23. Ex: Find the mass, in grams, of 3.0 mols of molecular oxygen, O2. 3.0 mol O2x32.00 g O2= 96 g O2 1 mol O2 96 g O2

  24. Ex: Find the number of moles in 236.5 g of CuSO4. 236.5 g CuSO4x1 mol CuSO4= 1.4817 mol CuSO4 159.61 g CuSO4 1.482 mol CuSO4

  25. Volume of a Mole of Gas • Standard temperature and pressure (STP) – conditions in which gas volumes are generally measured • Standard Temperature: 0 oC, 273 K, or 32 oF • Standard Pressure: 101.3 kPa, 1 atm, 760 mm Hg • Molar volume – 1 mol of any gas at STP takes up 22.4L of space.

  26. Ex: What is the volume of 0.960 mol of CH4 at STP? 0.960 mol CH4x22.4 L CH4= 21.504 L CH4 1 mol CH4 21.5 L CH4

  27. Ex: What is the volume of 1.5 mol of N2 at STP? 1.5 mol N2x22.4 L N2= 33.6 L N2 1 mol N2 34 L N2

  28. Ex: How many mols are in 2.50 L of CO2 at STP? 2.50 L CO2x1 mol CO2= 0.111607 mol CO2 22.4 L CO2 0.112 mol CO2

  29. Ex: What is the molar mass of a gas with a density of 1.964 g/L at STP? 1.964 gx22.4 L= 43.9936 g/mol L 1 mol 43.99 g/mol

  30. The Mole Road Map

  31. Ex: Calculate the number of molecules in 60.0 g NO2 60.0 g NO2 x 1 mol NO2 x 6.02 x 1023 molecules NO2 46.01 g NO2 1 mol NO2 7.85046729 x 1023 molecules 7.85 x 1023 molecules

  32. Ex: Calculate the volume, in liters, of 3.24 x 1022 molecules Cl2 at STP. 3.24 x 1022 molecules Cl2 x 1 mol x 22.4 L 6.02 x 1023 molecules 1 mol 1.21 L Cl2 at STP

  33. Section 7.2 Mole-Mass and Mole-Volume Relationships • Did We Meet Our Objectives? • Use the molar mass to convert between mass and moles of a substance • Use the mole to convert among measurements of mass, volume, and number of particles

  34. Section 7.3 Percent Composition and Chemical Formulas • Objectives: • Calculate the percent composition of a substance from its chemical formula or experimental data • Derive the empirical formula and molecular formula of a compound from experimental data

  35. Calculating the Percent Composition of a Compound • Percent composition – the relative amounts of each element in a compound • Percent by Mass

  36. Ex: An 8.20 g piece of magnesium combines completely with 5.40 g of oxygen to form a compound. What is the percent composition of this compound?

  37. Calculating the Percent Composition of a Compound • Percent composition – the relative amounts of each element in a compound • Percent by Composition

  38. Ex: Find the percent composition of propane (C3H8).

  39. Using Percent as a Conversion Factor • To do this, you multiply the mass of the compound by a conversion factor that is based on the percent composition. • Ex: Calculate the mass of carbon in 82.0 g of propane (C3H8). (Remember, carbon is 81.8%)

  40. Calculating Empirical Formulas • Empirical formula – gives the lowest whole number ratio of atoms of the elements in a compound • Empirical formula can sometimes be the molecular formula • CO2 • C6H12O6 CH2O

  41. Calculating Empirical Formulas The polymer used for the nonstick surface of cooking utensils is 24.0% C and 76.0% F by mass. • What is the empirical formula of this polymer?

  42. Calculating Molecular Formulas The polymer used for the nonstick surface of cooking utensils is 24.0% C and 76.0% F by mass. • What is the empirical formula of this polymer? • If the molecular mass is 100.02 g, what is the molecular formula?

  43. Section 7.3 Percent Composition and Chemical Formulas • Did We Meet Our Objectives? • Calculate the percent composition of a substance from its chemical formula or experimental data • Derive the empirical formula and molecular formula of a compound from experimental data

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