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Moles

Moles. Moles EQ: What is a mole? How do Scientist use the mole unit to count?. GPS: SC2 Students will relate how the Law of Conservation of Matter is used to determine chemical composition in compounds and chemical reactions.

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Moles

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  1. Moles

  2. MolesEQ: What is a mole?How do Scientist use the mole unit to count? • GPS: SC2Students will relate how the Law of Conservation of Matter is used to determine chemical composition in compounds and chemical reactions. c. Apply concepts of the mole and Avogadro’s number to conceptualize and calculate • Empirical/molecular formulas • Mass, moles and molecules relationships • Molar volumes of gases.

  3. MOLE • The chemistry term “mole” is used to describe a large amount. It makes counting groups of things easier. • SI base unit used to measure the amount of a substance. • Abbreviated: mol

  4. Mole • Mole • Avogadro’s number= 6.02 x 1023 • A mole of anything contains 6.02 x 1023 representative particles. • A particle can be anything – atoms, ions, electrons molecules, formula units, compound, etc.

  5. Molar Mass (mm) • Definition: mass in grams of one mole of any pure substance.Unit for molar mass: g/mol • mm = the atomic mass of that substance • Mn has atomic mass of 54.94 amu, on the Periodic Table; therefore its molar mass is 54.94 g/mol. • Ex. Find the molar mass of : Cu, Ba, Au, Ag, Hg.

  6. Molar Mass of Compounds • mm of Compounds= total mass of all atoms in compound. • It’s the sum of the mass of each atom of the element in the compound. • ∑ # of moles x molar mass of atom= # of grams 1. Ex. CaCl2 = 1 Ca + 2 Cl =40.08 g/mol + (2)35.5 g/mol = 40.08 + 71= 111 g/mol 2. Ex. CH3COOH 2 C x 12.01 g C = 24.02 g C + 4 H x 1.008 g H = 4.032 g H + 2 O x 16.00 g O = 32.00 g O mm of CH3COOH = 60.05 g CH3COOH

  7. Molar Mass of Compounds Practice Problems: Determine the molar mass of each: 1. MgCl2 8. AgNO3 2. KC2H3O2 9. CO2 3. (NH4)3PO4 10. Pb(NO3)2 4. Au 11. U 5. Ne 12. H2S2O8 6. NaCl 7. Ba(NO3)2

  8. Molar Conversions Mole x mm ÷ 6.02 x 1023 Mass Particle ÷ mm x 6.02 x 1023 Mole

  9. One-step Conversion:- • Use conversion chart. • Look at the starting unit, this is where you begin. This is your KNOWN • Determine what you are solving for, this is the ending unit. This is your UNKNOWN • Follow the arrow on the chart to convert. • Perform the correct operation, i.e. multiply or divide numbers.

  10. 2-step Conversion:same as before, just 2 parts- • Use conversion chart. • Look at the starting unit, this is where you begin. This is your KNOWN • Determine what you are solving for, this is the ending unit. This is your UNKNOWN • Follow the arrow on the chart to convert. • Perform the correct operation, i.e. multiply or divide numbers.

  11. Converting: Mol  Particles Particles  Mol 1) How many sucrose particles are there in 3.5 moles sucrose? 3.5 mol= _________particles sucrose Answer: 3.5 mol sucrose x 6.02 *1023 particles sucrose = 2.11 x 1024 particles sucrose 2) Determine the number of moles in 2.50 x 1018 atoms Zn. 2.50 x 1018 atoms Zn = ___________mol Zn Answer: 2.50 x 1018 atoms Zn / 6.02 *1023 atoms Zn = 4.153 x 10-6 mol Zn

  12. Moles  MassMass  Mole Example: • Calculate the mass (in g) of 0.0450 moles of Chromium. 0.0450 mol Cr= _________g Cr Answer: 0.0450 mol Cr x 52.00 g/mol Cr = 2.34 g Cr • How many moles of calcium are in 525 g Ca? 525 g Ca= _________ mol Ca Answer: 525 g Ca / 40.08 g/mol Ca = 13.1 mol Ca

  13. Moles → Mass of a Compound Mass of a Compound → Moles • Moles → Mass of a Compound • Calculate molar mass of compound first • mol known x mm • Mass of a Compound → Moles • Invert the conversion factor, divide. • Mass known / mm cmpd How many moles are in 325 g Ca(OH)2. Molar mass= 74.10 g/mol 325 g Ca(OH)2 / 74.10 g Ca(OH)2 = 4.39 mol Ca(OH)2

  14. Converting: Mol  Particles Particles  Mol Practice problems: • Determine the number of molecules in 35.3 mol CO2. 2. Convert 5.75 x 1024 atoms Al to moles. 3. Convert 3.75 x 1024 molecules CO2 to mol.

  15. Moles  MassMass  Mole • Practice Problems: • Determine the mass in grams of each: a. 3.57 mol Al b. 42.6 mol Si 2. Determine the number of moles in each: a. 300.0 g S b. 125 g Zn

  16. Moles → Mass of a Compound • Practice problem: • What is the mass of 4.35 x 10-2 moles of ZnCl2? • How many grams of potassium permanganate (KMnO4) are in 2.55 moles? • How many moles are in 35.5 g KMnO4?

  17. 2 step conversion:Mass  Mol  Particles • Mass  Mol  Particles formula: Known mass x 6.02 x 1023 rep. part. = # particles mm of substance • How many gold (Au) atoms are there in a 25.0 g gold nugget? 25.0 g Au= _______________ particles Au Answer: 25.0 g Au / 197 g/mol Au x 6.02 x 1023 atoms Au = 7.65 x 1022 atoms Au

  18. 2 step conversion:Particles  Mol  Mass 2. Particles  Mol  Mass formula: Known particles x mm of substance = mass (g) 6.02 x 1023 • What is the mass in grams of 5.50 x 1022 atoms helium? 5.50 x 1022 atoms He= __________ g He Answer: 5.50 x 1022 atoms Hex 4.00 g/mol He 6.02 x 1023 = 0.33 g He

  19. Practice How many particles are in each of the following: • 55.2 g LiCl • 0.230 g Pb(NO3)2 • 11.5 g Hg What is the mass in grams of each? • 6.02 x 1024 atoms Bi • 1.00 x 1024 particles MnO2 • 3.40 x 1022 atoms He • 1.50 x 1015 molecules NH3

  20. Percent Composition(% by mass) • Percent composition (% by mass): The percent by mass of each element in a compound. • mass of element x 100 = % Composition mass of cmpd Example: Determine the % composition of H2O: 2.02 g H x 100= 11.2% H 16.00 g O x 100 = 88.8% O 18.02 g H20 18.02 g H20

  21. Percent Composition(% by mass) • Practice Problem: • Determine the % by mass of each element in: a. CaCl2. b. Na2SO4. 2. Which has the larger % by mass of sulfur, H2SO3 or H2S2O8?

  22. Empirical formula: formula with the smallest whole-number mole ratio of the elements. Molecular formula: actual number of atoms of each element in the compound Empirical & Molecular Formulas:*subscripts= ratio of each element in compound*

  23. Empirical Formula (EF) FOLLOW THESE STEPS: 1. Calculate the % by mass. Assume %=100 g, just change % sign to g. 2. Convert known mass from step 1 to moles (mass→mol). Go to step 6 if all #’s are whole #’s. 3. Write ratio to compare the # of moles of each atom in the cmpd. 4. Find simplest whole # ratio by dividing each subscript/ ratio by the smallest subscript to get whole #’s. 5. (If still no whole #: multiply each subscript by the smallest whole # that will convert all subscripts to a whole # or round to a whole #). 6. Write Empirical Formula. (Symbols + subscript)

  24. EF Rhyme • % to mass • Mass to mole • Divide by small • Multiply til whole

  25. Molecular Formula (MF) For Molecular Formula: 1st do steps 1-6 for EF. 7. Determine (n)= whole # factor that the subscripts of the E.F. are multiplied. To get n: n = experimental mass (given mass) mass of EF • MF= n(EF) multiply n by each subscript in EF to get MF.

  26. EF practice • Calculate the EF of a cmpd that contains 38.71% C, 9.68 % H, & 51.61% O. • Analysis of a cmpd indicates the % composition 42.07% Na, 18.89% P, & 39.04% O. Determine its EF. • Analysis of a chemical used in photographic developing fluid indicates a chemical composition of 65.45% C, 5.45% H, & 29.09% O. Determine the EF.

  27. Empirical & Molecular Formulas • Practice problems: • Analysis of a chemical used in photographic developing fluid indicates a chemical composition of 65.45% C, 5.45% H, and 29.09% O. The molar mass is found to be 110.0g/ mol. Determine the empirical and molecular formulas. • A compound was found to contain 49.98 g carbon and 10.47 g hydrogen. The molar mass of the compound is 58.12 g/mol. Determine the E.F and M.F.

  28. Empirical & Molecular Formulas 3. A colorless liquid composed of 46.68% N and 53.32% O has a molar mass of 60.01 g/ mol. What are the empirical & molecular formulas?

  29. The END

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