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“THE MOLE”

“THE MOLE”.

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“THE MOLE”

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  1. “THE MOLE”

  2. Counting Quantities1.) You can count how many of something you have……12 doughnuts in a dozen10 pins knocked down for a strike2.) You can measure the mass of what you have…….buying candy by the poundgold is priced by the gram3.) You can measure matter by volume…..gasoline by the litermedicine in mL

  3. “THE MOLE” (Molar Mass = 1 MOLE) A convenient standard quantity of a substance For solids, liquids, gases…..you can convert weight to moles, or moles to weight The amount of a substance = in grams to the sum of the atomic weights

  4. 1 mole =6.02 x 1023 particles • What do I mean when I say particles? • Can be atoms, molecules, formula units, or ions. ELEMENT use atoms COVALENT COMPOUND use molecules IONIC COMPOUND use formula units ION use ions

  5. Examples • Na • I mole of Na = 6.02 x 1023 atoms of Na • H2O • 1 mole of H2O = 6.02 x 1023 molecules of H2O • NaCl • 1 mole of NaCl = 6.02 x 1023 formula units of NaCl • Cl- • 1 mole of Cl- = 6.02 x 1023 ions of Cl-

  6. MOLAR MASS AND NUMBER OF PARTICLES SubstanceMolar MassNumber of Particles in One Mole • Carbon (C) 12.0 g 6.02 x 1023 C atoms • Sodium (Na) 23.0 g 6.02 x 1023 Na atoms • CaCO3 (antacid) 100.1 g 6.02 x 1023 CaCO3 formula units • C6H12O6 (glucose) 180.0 g 6.02 x 1023 glucose molecules • C8H10N4O2 (caffeine) 194.0 g 6.02 x 1023 caffeine molecules

  7. Example problems • How many moles is 3.46 x 1026 molecules of water? • How many moles of hydrogen are in 3.46 x 1026 molecules of water? • How many atoms is 3.46 moles of silver?

  8. A dozen… • Will a dozen trucks have the same mass as a dozen chocolates? • Will 6.02 x 1023 molecules of water have the same mass as 6.02 x 1023 molecules of CO2?

  9. Atomic Mass Unit (AMU) • AMU = mass of a single atom • AMU are relative values based on the mass of the most common isotope of carbon.

  10. Molar Mass • More convenient to use UNITS: g/mol • MM= mass of one mole of a substance….remember???

  11. Molar Mass Use periodic table to find MM of a compound: 1. Multiply each element’s mass (from the periodic table) by number of atoms present 2. Add all the masses together

  12. Practice Problem • Find the molar mass of (NH4)2O Did you get 52 g/mol? • How many grams is 2.35 moles of MgO?

  13. Practice Problem • How does the amount of moles of 44 g of CO2 compare to 36.6 g of HCl? • Same amount of moles; therefore, same amount of molecules (6.02 x 1023)

  14. Example • How many moles is 300 g CO2? • 1. Determine the relationship between g CO2 and moles of CO2 (find molar mass) • 2. Cancel out units and solve.

  15. Solution 44 g CO2 = 1 mole CO2 300 g CO2 1 mole CO2 44 g CO2 = 6.82 mole CO2

  16. Example • How many molecules is 6.82 moles CO2? • Determine the relationship between moles of CO2 and molecules CO2 (Avogadro’s number) • 2. Cancel out units and solve.

  17. Solution 6.82 mole CO2 6.02 x 1023 molecules 1 mole CO2 = 4.11 x 1024 molecules CO2

  18. NOTICE: to get from grams to molecules you have to go through the MOLE • Remember in chemistry all roads lead to the mole!

  19. MOLAR VOLUME • PV=nRT • What is the volume of 1 mole of gas of oxygen at STP? • What is the volume of 1 mole of gas of hydrogen at STP?

  20. MOLAR VOLUME • Gases can change volume based on temperature and pressure • BUT at STP 6.02 x 1023 molecules (1 mole) = 22.4 L • This is true of ANY gas! • How many liters are in 12.3 moles of H2SO4 (gas) at STP?

  21. 12.3 moles H2SO4 22.4 Liters 1 mole H2SO4 = 276 liters H2SO4

  22. Mole Road Map

  23. Example • What volume is occupied by 90 g CO2 at STP? • 1. Determine conversion factors • 2. Use factor label

  24. 1 mole CO2 = 44 g CO2 1 mole CO2 (at STP) = 22.4 L CO2 1 mole CO2 22. 4 L CO2 90 g CO2 1 mole CO2 44 g CO2 = 45.8 L CO2

  25. HAPPY VALENTINE’S DAY • Make a poem!!! ROSES ARE RED VIOLETS ARE BLUE MOLAR MASS SURE IS SWEET AND MOLAR VOLUME IS TOO

  26. MOLE MOVIE • http://www.learner.org/resources/series61.html#

  27. BELLWORK 2-16 • Convert 3.49 x 1029 ions of PbCl2 to moles. • What is the mass of hydrogen in 50.0 grams of C3H8?

  28. PERCENT COMPOSITION • Relative amounts of elements in a compound are expressed as PERCENT COMPOSITION.

  29. % composition steps • Step 1: Find the total mass of the compound • Step 2: Divide the total mass of each element by the total mass of the compound • Step 3: Multiply by 100

  30. Example • K2CrO4 • Total mass = 2 K (39.1 x 2) + Cr (52) + 4 O (16 x 4) 194.2 g • K: 78.2 g/ 194.2 g = .403 x 100 = 40.3% K

  31. Cr: 52g/ 194.2 g = .268 x 100 = 26.8 % Cr • O: 64 g/ 194.2 g = .329 x 100 = 32.9 % O 32.9 % + 26.8 % + 40.3% = 100%

  32. CAFFEINE: C8H10N4O2

  33. EMPIRICAL FORMULAS • The lowest whole-number ratio of the atoms of the elements in a compound • Is C2O4 an empirical formula? • CO2 1:2 ratio between Carbon and Oxygen

  34. Is H2SO4 an empirical formula? • Yes, because you can’t reduce the subscripts any further. (2:1: 4 ratio)

  35. A compound is analyzed and found to contain 25.9% nitrogen and 74.1% oxygen. What is the empirical formula of the compound?

  36. STEPS FOR EMPIRICAL FORMULA PROBLEMS • sTeP 1: change all percents of elements to grams • sTeP 2: divide grams of each element by molar mass of each element- this will give you moles of each element • sTep 3: divide moles of each element by the smallest amount of moles • sTeP 4: if the ratios are not in whole numbers multiply each by the smallest whole number that will convert both subscripts to whole numbers

  37. Example Problem • Step 1: 25.9 grams of Nitrogen 74.1 grams of Oxygen • Step 2: 25.9 g x 1 mol = 1.85 mol N 14 g 74.1 g x 1mol = 4.63 mol O 16 g • Step 3: 1.85 mol N/ 1.85 = 1 mol N 4.63 mol O/ 1.85 = 2.5 mol O • Step 4: multiply mol N and mol O by 2 to get whole numbers: 2 mol N and 5 mol O • Empirical formula equals N2O5

  38. Class work • What is the empirical formula of a compound that is 52.5% Phosphorus and 47.5% Oxygen? • What is the empirical formula of a compound that is 10.0% C, .80% H, and 89.1% Cl? • What is the empirical formula of a compound that is 7.8% C and 92.2% Cl?

  39. Did you get these answers? • P4O7 • CHCl3 • CCl4

  40. Bellwork 2/20 • #1: Which of the following can be classified as an empirical formula? a) SCl2 b) C6H10O4 c) Na2SO3 • #2: What is the empirical formula of a compound that is 60.0 % Magnesium and 40.0% oxygen?

  41. SOLUTION • #1: a and c are empirical formulas • #2: Change % to grams • 60.0% Magnesium= 60.0 g Magnesium • 40.0 % Oxygen = 40.0 g Oxygen

  42. 2. Convert the grams of each element to moles (using the molar mass) 60.0 g Magnesium 1 mole Magnesium 24.3 g Magnesium = 2.46 moles Magnesium

  43. Solution Continued • Oxygen: 40.0 g Oxygen 1 mole Oxygen 16.0 g Oxygen = 2.50 moles Oxygen

  44. Solution Continued • 3. Now divide the number of moles for each element by the smallest number of moles. • 2.46 moles Magnesium • 2.50 moles Oxygen • Magnesium has the smallest number of moles {2.46 moles} • So divide the number of moles of Magnesium and Oxygen by 2.46

  45. Solution Continued • Magnesium = 2.46 moles / 2.46 moles = 1 Mg • Oxygen = 2.50 moles / 2.46 moles = 1.02 O • Mg1O1 • MgOis the empirical formula

  46. MOLECULAR FORMULAS • Is a simple whole-number multiple of the empirical formula • Can be the same as its experimentally determined empirical formula. • EXAMPLE: C8H16 is the molecular formula CH2 would be the empirical formula

  47. STEPS FOR FINDING MOLECULAR FORMULA • Step 1: First find the empirical formula following the steps given previously (sometimes it may be given to you) • Step 2: Divide molar mass of compound by empirical formula mass. • Step 3: Multiply each subscript in the empirical formula by the value you got in step 2.

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