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Chapter 10 “Chemical Quantities”

Chapter 10 “Chemical Quantities”. Yes , you will need a calculator for this chapter!. Section 10.1 The Mole: A Measurement of Matter. OBJECTIVES: Describe methods of measuring the amount of something. Define Avogadro’s number as it relates to a mole of a substance.

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Chapter 10 “Chemical Quantities”

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  1. Chapter 10“Chemical Quantities” Yes, you will need a calculator for this chapter!

  2. Section 10.1The Mole: A Measurement of Matter • OBJECTIVES: • Describe methods of measuring the amount of something. • DefineAvogadro’s number as it relates to a mole of a substance. • Distinguish between the atomic mass of an element and its molar mass. • Describe how the mass of a mole of a compound is calculated.

  3. How do we measure items? • We measure mass in grams. • We measure volume in liters. • We count pieces in MOLES.

  4. What is the mole? We’re not talking about this kind of mole!

  5. Moles (is abbreviated: mol) • It is an amount, defined as the number of carbon atoms in exactly 12 grams of carbon-12. • 1 mole = 6.02 x 1023 of the representative particles. • Treat it like a very large dozen • 6.02 x 1023 is called: Avogadro’s number.

  6. Similar Words for an amount • Pair: 1 pair of shoelaces = 2 shoelaces • Dozen: 1 dozen oranges = 12 oranges • Gross: 1 gross of pencils = 144 pencils • Ream: 1 ream of paper = 500 sheets of paper

  7. What are Representative Particles? • The smallest pieces of a substance: • For a molecular compound: it is the molecule. • For an ionic compound: it is the formula unit (made of ions). • For an element: it is the atom. • Remember the 7 diatomic elements? (made of molecules)

  8. Types of questions • How many oxygen atoms in the following? CaCO3 Al2(SO4)3 • How many ions in the following? CaCl2 NaOH Al2(SO4)3 3 atoms of oxygen 12 (3 x 4) atoms of oxygen 3 total ions (1 Ca2+ ion and 2 Cl1- ions) 2 total ions (1 Na1+ ion and 1 OH1- ion) 5 total ions (2 Al3+ + 3 SO42- ions)

  9. Measuring Moles • Remember relative atomic mass? - The amu was one twelfth the mass of a carbon-12 atom. • Since the mole is the number of atoms in 12 grams of carbon-12, • the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.

  10. Gram Atomic Mass (gam) • Equals the mass of 1 mole of an element in grams (from periodic table) • 12.01 grams of C has the same number of pieces as 1.01 grams of H and 55.85 grams of Fe. • We can write this as: • 12.01 g C = 1 mole C (this is also the molar mass) (**Remember we can count things by weighing them.)

  11. nitrogen aluminum zinc Find the Gram Atomic Mass (gam) of the following: 14.01 g 26.98 g 65.39 g

  12. What about compounds? In 1 mole of H2O molecules there are 2 moles of H atoms and 1 mole of O atoms (think of a compound as a molar ratio) • To find the mass of one mole of a compound: • Determine the number of moles of the elements present • Multiply the number times their mass (from the periodic table) • Add them up for the total mass

  13. Calculating Gram Formula Mass (gfm) Calculate the formula mass of magnesium carbonate, MgCO3. 84.32 g 24.31 g + 12.01 g + 3 x (16.00 g)= Thus, 84.32 grams is the formula mass for MgCO3.

  14. GFM Practice Problem: What is the mass of one mole of CH4? 1 mole of C = 12.01 g 4 mole of H x 1.01 g = 4.04g 1 mole CH4 = 12.01 + 4.04 = 16.05g

  15. Section 10.2Mole-Mass and Mole-Volume Relationships • OBJECTIVES: • Describe how to convert the mass of a substance to the number of moles of a substance, and moles to mass. • Identify the volume of a quantity of gas at STP.

  16. Molar Mass • Molar mass is the generic term for the mass of one mole of any substance (expressed in grams/mol) • The same as: 1) Gram Molecular Mass (for molecules) 2) Gram Formula Mass (ionic compounds) 3) Gram Atomic Mass (for elements) • molar mass is just a much broader term than these other specific masses

  17. Examples • Calculate the molar mass of the following and tell what type it is: Na2S N2O4 C Ca(NO3)2 C6H12O6 (NH4)3PO4 = 78.05 g/mol (gram formula mass) = 92.02 g/mol (gram molecular mass) = 12.01 g/mol (gram atomic mass) = 164.10 g/mol (gram formula mass) = 180.18 g/mol (gram molecular mass) = 149.12 g/mol (gram formula mass)

  18. Since Molar Mass is… • The number of grams in 1 mole of atoms, ions, or molecules • We can make conversion factors from these. - To change between grams of a compound and moles of a compound.

  19. For example How many moles is 5.69 g of NaOH?

  20. For example How many moles is 5.69 g of NaOH?

  21. For example How many moles is 5.69 g of NaOH? • need to change grams to moles

  22. For example How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH

  23. For example How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g

  24. For example How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  25. For example How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  26. For example How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

  27. Practice Problems: • How much would 2.34 moles of carbon weigh? • How many moles of magnesium is 24.31 g of Mg? 28.1 grams C 1.000 mol Mg

  28. Avog. # Practice problems: • How many molecules of CO2 are in 4.6 moles of CO2? • How many moles of water is 5.87 x 1022 molecules? • How many atoms of carbon are in 1.230 moles of Carbon? • How many moles is 7.78 x 1024 formula units of MgCl2? 2.8 x 1024 m.c. CO2 0.0975 mol H2O (or 9.75 x 10-2) 7.405 x 1023 atoms C 12.9 moles MgCl2

  29. The Mole-Volume Relationship • Many of the chemicals we deal with are: gases. • Two things effect the volume of a gas: a) Temperatureand b) Pressure **We need to compare all gases at the same temperature and pressure.

  30. Standard Temperature and Pressure abbreviated STP 0ºC (273K) and 1.0 atm pressure At STP 1 mole of gas occupies 22.4 L= molar volume 1 mole = 22.4 L of any gas at STP

  31. Practice Problems: • What is the volume of 4.59 mole of CO2 gas at STP? • How many moles is 5.67 L of O2at STP? = 103 L = 0.253 mol

  32. Summary: These four items are all equal: a) 1 mole b) molar mass (in grams/mol) c) 6.02 x 1023 particles (atoms, molecules, or formula units) d) 22.4 L of a gas at STP **Thus, we can make conversion factors from them.

  33. Mixed Practice Problems: • How many atoms of lithium is 1.0 g of Li? • How much would 3.45 x 1022 atoms of U weigh? • What is the volume of 8.8 g of CH4 gas at STP? 8.7 x 1022 atoms Li 13.6 grams U = 12 L

  34. Section 10.3Percent Composition and Chemical Formulas • OBJECTIVES: • Describe how to calculate the percent by mass of an element in a compound. • Interpret an empirical formula. • Distinguish between empirical and molecular formulas.

  35. Percentage Composition the percentage by mass of each element in a compound

  36. Calculating Percent Composition of a Compound • Like all percent problems: part whole • Find the mass of each of the components (the elements masses from the periodic table ), • Next, divide by the total mass of the compound; then x 100 x 100 % = percent

  37. Percentage Composition Find the % composition of Cu2S. 127.10 g Cu 159.17 g Cu2S 32.07 g S 159.17 g Cu2S  100 = %Cu = 79.85% Cu %S =  100 = 20.15% S

  38. Find the mass percentage of water in calcium chloride dihydrate, CaCl2•2H2O? Percentage Composition 36.04 g 147.02 g 24.51% H2O %H2O =  100 =

  39. Empirical Formula (EF) Lowest whole number ratio of atoms in a compound CH3 C2H6 reduce subscripts

  40. Empirical Formula (EF) • Just find the lowest whole number ratio C6H12O6 CH4N • A formula is not just the ratio of atoms, it is also the ratio of moles. • In 1 mole of CO2 there is 1 mole of carbon and 2 moles of oxygen. • In one molecule of CO2 there is 1 atom of C and 2 atoms of O. = CH2O = this is already the lowest ratio.

  41. Formulas(continued) Formulas for ionic compounds are ALWAYS empirical (the lowest whole number ratio = cannot be reduced). Examples: NaCl MgCl2 Al2(SO4)3 K2CO3

  42. Formulas(continued) Formulas for molecular compoundsMIGHT be empirical (lowest whole number ratio). Molecular: H2O C6H12O6 C12H22O11 (Correct formula) Empirical: H2O CH2O C12H22O11 (Lowest whole number ratio)

  43. Calculating EF We can get a ratio from the percent composition. 1.) Assume you have 100 g. -the percentage becomes grams (75.1% = 75.1 grams) 2.) Convert grams to moles. 3.) Find lowest whole number ratio by dividingby the smallest# of moles. 4.) If not a whole #, use a multiplier.

  44. Example Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N. Assume 100 g so: 38.67 g C x 1mol C = 3.220 mole C 12.01 gC 16.22 g H x 1mol H = 16.06 mole H 1.01 gH 45.11 g N x 1mol N = 3.220 mole N 14.01 gN Now divide each value by the smallest value

  45. Example 3.220 mol C = 1 mol C 3.220 16.06 mol H = 5 mol H 3.220 3.220 mol N = 1 mol N 3.220 EF = C1H5N1 = CH5N

  46. Empirical Formula (EF) Find the empirical formula for a sample of 25.9% N and 74.1% O. 1.85 mol 1.85 mol 25.9 g 1 mol 14.01 g = 1.85 mol N = 1 N 74.1 g 1 mol 16.00 g = 4.63 mol O = 2.5 O

  47. Empirical Formula N1O2.5 N2O5 Need to make the subscripts whole numbers  multiply by 2

  48. Molecular Formula (MF) “True Formula” - the actual number of atoms in a compound C2H6 CH3 empirical formula ? molecular formula

  49. Empirical to Molecular Since the empirical formula is the lowest ratio, the actual molecule would weigh more. By a whole number multiple. **Divide the actual molar mass by the empirical formula mass.

  50. Molecular Formula The empirical formula for ethylene is CH2. Find the molecular formula if the molecular mass is 28.1 g/mol? 28.1 g/mol 14.03 g/mol empirical mass = 14.03 g/mol = 2.00 (CH2)2 C2H4

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