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Mastering the Mole: A Comprehensive Guide to Chemical Quantities

Dive into the world of chemical quantities with this detailed guide on the mole, Avogadro’s number, and mass calculations for molecular and ionic compounds. Understand conversion factors, molar mass, and volume relationships for gases at STP. Learn to calculate molecules, volumes, and percent composition, and derive chemical formulas from experimental data. This resource-rich chapter covers essential concepts and provides practice problems for comprehensive learning.

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Mastering the Mole: A Comprehensive Guide to 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. 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. Molecular compounds Ionic compounds

  4. 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

  5. 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 particles • 1 mole = 6.02 x 1023 ions

  6. 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

  7. Ex: How many moles of magnesium is 1.25 x 1023 atoms of magnesium?

  8. 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

  9. Ex: How many molecules is 0.360 mol of water?

  10. 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

  11. Ex: How many atoms are there in 1.14 mol SO3?

  12. Ex: How many atoms are in 2.12 mol of propane (C3H8)?

  13. 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.

  14. 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

  15. 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 molecular mass of the following: • CaI2 • (NH4)2CO3

  16. 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

  17. 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

  18. 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

  19. Find the mass, in grams, of 2.5 mols of Na. • Find the number of mols in 75.0 g of dinitrogen trioxide (N2O3).

  20. Find the mass, in grams, of 3.0 mols of molecular oxygen • Find the number of moles in 236.5g of CuSO4

  21. 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.

  22. Volume of a Mole of Gas • What is the volume of 0.960 mol of CH4 at STP? • What is the volume of 1.5 mol of N2 at STP?

  23. Volume of a Mole of Gas • How many mols are in 2.50L of CO2 at STP? • What is the molar mass of a gas with a density of 1.964 g/L?

  24. The Mole Road Map

  25. Calculate the number of molecules in 60.0 g NO2 • Calculate the volume, in liters, of 3.24 x 1022 molecules Cl2 at STP.

  26. 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

  27. 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

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

  29. 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?

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

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

  32. 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%)

  33. 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

  34. Calculating Molecular Formulas • Show on board & examples

  35. 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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