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Chapter 10

Chapter 10. Chemical Quantities. Molar Mass. MOLAR MASS—# of grams of an element in one mole of that element. = atomic mass of an element with the units of g/mol Get it from the periodic table. Molar Mass. Elements. Find the molar mass on the periodic table round to 0.1 g/mol.

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Chapter 10

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  1. Chapter 10 Chemical Quantities

  2. Molar Mass • MOLAR MASS—# of grams of an element in one mole of that element. = atomic mass of an element with the units of g/mol • Get it from the periodic table

  3. Molar Mass Elements • Find the molar mass on the periodic table round to 0.1 g/mol Compounds • ADD the molar mass of every atom present in the compound

  4. Molar Mass Practice

  5. If you know the MOLAR MASS … • You can convert between grams and moles! Volume of gas at STP 1 Mole Molar mass (g) Mass Representative Particles

  6. Mole ↔ Mass Examples • How many moles in 51.0 g of Na3PO4? • Find Molar Mass: (23.0 x 3) +31.0+(16.0 x 4) = 164.0 g/mol • 51.0 g Na3PO4 x 1 mole Na3PO4= 0.311 mol Na3PO4 164 g Na3PO4 • What is the mass of 0.70 mol of NH4Cl? • Find Molar Mass: 14.0+(1.0 x 4) +35.5 = 53.5g/mol • 0.70 mol NH4Cl x 53.5 g NH4Cl = 37.45 g NH4Cl 1 mol NH4Cl

  7. 1 mole= Avogadro’s # = 6.02x1023 representative particles • How big is 6.02x1023? 602 000 000 000 000 000 000 000 • A representative particle: atom, molecule, or formula unit • 1 mole of nitrogen gas (N2) contains 6.02 x 1023molecules. • Because N2 is a molecule (Covalent) • 1 mole of calcium fluoride (CaF2) contains 6.02 x 1023formula units • Because CaF2 is a formula unit (ionic) • How many atoms of He are in 1 mole?

  8. Volume of gas at STP Avogadro’s Number 6.02 x 10 23 1 Mole Molar mass Mass Representative Particles Unit: Grams (g) Type of Substance:Unit: Elements Atoms ionic Formula Units (f.un) Covalent Molecules (mocs)

  9. By COUNT: Avogadro's NumberHow big is 6.02x1023?602 000 000 000 000 000 000 000 • If you had Avogadro's number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. • If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole. • A mole of coke cans would cover the surface of the earth to a depth of over 200 miles.

  10. Measuring Matter • Some units of measurement indicate specific numbers. • A pair means 2 • A dozen means 12 • Knowing how count, mass, & volume relate allows you to convert between them. • If 1 dozen apples = 12 apples, and 1 dozen apples has a mass of 2.0 kg, What is the mass of 90 apples? • 90/12 = 7.5 dozen 7.5 dozen x 2.0 kg • = 15 kg

  11. What is a Mole? • When dealing with tiny particles (atoms, ions, compounds), the sample size is usually very large. • Counting is not practical. • Just as a dozen eggs represents 12 eggs, a mole (mol) of a substance represents 6.02 x 1023 representative particles of that substance. • The number 6.02 x 1023 is known as Avogadro’s number

  12. Moles • Determining the number of atoms in a mole of a compound: • How many moles are in a representative particle (formula unit or molecule) of the substance? • This can be determined from the formula. • Example: each molecule of CO2 contains 3 atoms: 1 carbon atom and 2 oxygen atoms. • A mole of CO2 contains 1 mole of carbon atoms and 2 moles of oxygen atoms.

  13. The Mass of a Mole • When dealing with atoms, it is often easier to work with mass. • Gram atomic mass (gam)—the atomic mass of an element expressed in grams. • The atomic mass of carbon is 12 amu • The gam of carbon is 12 g. • The gam is the mass of 1 mole of atoms of any element.

  14. The Mass of a Mole of a Compound • To determine the mass of a mole of a compound you need to know the formula and the gam of each atom in the compound. • Add the masses of each atom to get the mass of the compound • Sulfur trioxide (SO3) • 1 mole of S (32.1g) & 3 moles of O (16.0 g each) • 32.1g + 16.0g + 16.0g + 16.0g = 80.1 g • Gram molecular mass (gmm)—the mass of 1 mole of a molecular compound • Gram formula mass (gfm)—the mass of 1 mole of an ionic compound

  15. Sec. 2 Mole-Mass-Volume • We learned about gam, gmm, & gfm last time. We can use 1 broad term to tell the mass of a substance. • Molar Mass—the mass (in grams) of 1 mole of a substance. • Why do we have the 3 terms then? Sometimes the term molar mass in unclear. What is the molar mass of oxygen? Do you mean oxygen gas (O2)? Then the molar mass is 32.0g (2 x 16.0g). Or do you mean oxygen atoms (O)? Then the molar mass is 16.0 g.

  16. The Volume of a Mole of Gas • The volumes of 1 mole of different solid and liquid substances are not the same. • The volumes of 1 mole of different gases are the same under the same conditions. • To keep things under the same conditions, gases are measured at standard temperature and pressure (STP) • Standard temp is 0°C • Standard pressure is 101.3kPa (or 1 atm) • At STP, 1 mole of any gas has a volume of 22.4 L

  17. Converting Between moles, particles, mass, and volume Volume of gas at STP Note: to convert between particles, mass, and volume, you have to go through moles. 1 mol 22.4 L 22.4 L 1 mol Mole 6.02 x 1023 particles 1 mol Molar mass 1 mol 1 mol . molar mass 1 mol . 6.02 x 1023 particles Mass Representative Particles

  18. Sec. 3: Percent Composition & Chemical Formulas • Percent Composition—The relative amount of each element in a compound • The % of all elements in the compound must equal 100% • % mass of element = grams of element x 100 grams of compound • Or • % mass of element = molar mass of element x 100 molar mass of compound

  19. Example • An 8.20 g piece of magnesium combines completely with 5.40 g of oxygen to form a compound. What is the % composition? • First—add 8.20 g & 5.40 g to get the mass of the compound. 8.20 + 5.40 = 13.60 g • % Mg = (mass of Mg/mass of compound) x 100 • % Mg = 8.20g/13.60g x 100 = 60.3% • % O = (mass of O / mass of compound) x 100 • % O = 5.40/13.60 x 100 = 39.7% • Does this make sense? • 60.3 + 39.7 = 100

  20. Empirical Formulas • Empirical formula—gives the lowest whole-number ratio of the elements in a compound • An empirical formula may or may not be the same as a molecular formula. • If the formulas are different, the molecular formula is a simple multiple of the empirical formula. • Examples: • The empirical formula for H2O2 is HO • For CO2, the empirical & molecular formula are the same. • C6H6 and C2H2 have the same empirical formula: CH

  21. Molecular Formulas • You can determine the molecular formula if you know empirical formula and molar mass. • Divide the molar mass by the empirical formula mass. • Multiply this number by all subscripts in the empirical formula to get the molecular formula.

  22. Example: • Find molecular formula with a molar mass of 60.0g and empirical formula of CH4N • 1st find the empirical formula mass • 1 C, 4 H, 1 N • 12 + (4 x 1) + 14 = 30 g • Then, divide molar mass by empirical mass • 60.0g / 30 g = 2 • Multiply each element subscript by the this. • C2H8N2

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