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Chapter 10 “Chemical Quantities”

Chapter 10 “Chemical Quantities”. Y ou will need a calculator for this chapter!. Section 10.1 p. 287 The Mole: A M easurement of Matter. How do we measure items?. You can measure mass , volume , or count pieces We measure mass in grams We measure volume in liters

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Chapter 10 “Chemical Quantities”

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

  2. Section 10.1 p. 287The Mole: A Measurement of Matter

  3. How do we measure items? • You can measuremass, • volume, • orcount pieces • We measure mass in grams • We measure volume in liters • We count pieces inMOLES

  4. Other Ways to Measure Amount • Pair: 1 pair of socks = 2 socks • Dozen: 1 dozen donuts = 12 donuts • Gross: 1 gross of pencils = 144 pencils (12 dozen) • Ream: 1 ream of paper = 500 sheets of paper Guided Practice Problem p. 289

  5. Practice Problem #2 pg. 289 • Assume 2.0 kg of apples is 1 dozen and that each apple has 8 seeds. How many apple seeds are in 14 kg of apples?(work INDEPENDENTLY to solve)

  6. What is the mole? Not this kind of mole!

  7. Moles (abbreviated mol) • Derived from German word molekül (molecule) • SI measurement of an amount • 1 mole = 6.02 x 1023of representative particles, or….. • # of carbon atoms in exactly 12 g of Carbon-12 isotope • Called Avogadro’s number

  8. What are Representative Particles?(Table 10.1 p. 290) • The smallest pieces of a substance: • molecular cmpd -molecule • ionic cmpd -formula unit(made of ions) • element: is the • Remember the 7 diatomic elements? (made of molecules) BrINClHOF atom Br2 I2 N2 Cl2 H2 O2 F2 Guided Practice Problem #3 p. 291

  9. Mole Video 3:49

  10. Quick Quiz 6.02 x 1023 • How big is a mole? • If everyone in the world got a mole of pennies, how much $ would every person have? • If you stacked a mole of paper how many times would it go from the Earth to the moon? • How long would it take for every person in the world to eat through a mole of marshmellows? 1 trillion bucks $1,000,000,000,000 80 billion times 80,000,000,000 40,000,000 years w/o a bathroom break!

  11. Consider these 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)

  12. Practice problems

  13. The Mass of a Mole of an Element • Atomic mass of element (mass of 1 atom) expressed in amu • atomic masses - relative masses based on mass of C-12 (12.0 amu) • 1 amu is 1/12 mass of C-12 atom

  14. Molar Mass…. • = mass of 1 mol of elementin grams(periodic table) • 12.01 grams C has same # particles as 1.01 g H & 55.85 g Fe • 12.01 g C = 1 mol C • 1.01 g H = 1 mol H • 55.85 g Fe = 1 mol Fe All contain 6.02 x 1023 atoms

  15. Molar Mass Practice Problems

  16. What about compounds? • 1 mol of H2O molecules has 2 mol of H atoms & 1 mol of O atoms (think of a compound as amolar ratio) • To find mass of 1 mol of a cmpd: • determine # moles of elements present • Multiply # times their mass (from periodic table) • add up for total mass

  17. Calculating Molar Mass Calculate molar mass of magnesium carbonate, MgCO3. 84.3 g 24.3 g + 12.0 g + 3 x (16.00 g) = So, 84.3 g = molar mass for MgCO3

  18. Section 10.2 p. 297Mole-Mass and Mole-Volume Relationships

  19. Molar Mass • Molar mass - generic term for mass of 1 mol ofanysubstance (expressed ingrams/mol) • Same as: 1) Gram Molecular Mass (for molecules) 2) Gram Formula Mass (ionic compounds) 3) Gram Atomic Mass (for elements) • molar mass is more broad term than these other specific masses

  20. Examples • Calculate the molar mass of: Na2S N2O4 C Ca(NO3)2 C6H12O6 (NH4)3PO4 = 78.05 g/mol = 92.02 g/mol = 12.01 g/mol = 164.10 g/mol = 180.12 g/mol = 149.12 g/mol

  21. Molar Mass is… • # of g in 1 mol of atoms, formula units, or molecules • Make conversion factors from these - To change btwn g of cmpd and mol of cmpd

  22. Using the Mole Roadmap • How many moles is 5.69 g of NaOH? 0.142 mol NaOH

  23. The Mole-Volume Relationship • gases - hard to determine mass • how many moles of gas? • 2 things affect gas V: • a)Temp& b)Pressure • compare all gases at = temp & pressure

  24. Standard Temperature and Pressure • 0ºC & 1 atm pressure - abbreviated “STP” • At STP, 1 mol ofanygas has V of 22.4 L - Calledmolar volume • 1 mol of any gas at STP = 22.4 L

  25. Practice Examples

  26. Mole Day Celebrated on October 23rd from 6:02 am until 6:02 pm (6:02 on 10-23)

  27. Density of a gas • D = m / V (density = mass/volume) - for gas units are: g / L • find density of a gas at STP if formula known • You need: 1) mass and 2) volume • Assume 1 mol, so massismolar mass (from periodic table) • At STP, V = 22.4 L

  28. Practice Examples (D=m/V)

  29. Another way: • If given density, find molar mass of gas • Assume 1 mol at STP, so V = 22.4 L modify: D = m/V to show: • “m” will be mass of 1 mol, given 22.4 L • What is molar mass of a gas with density of 1.964 g/L? • How about a density of 2.86 g/L? m = D x V = 44.0 g/mol 64.0 g/mol

  30. Summary • all equal: a) 1 mole b) molar mass (in grams/mol) c) 6.02 x 1023 representative particles (atoms, molecules, or formula units) d) 22.4 L of gas at STP makeconversion factorsfrom these 4 values (p.303)

  31. Copy this conversion map into your notes! Notice all conversions must go through the MOLE!

  32. Section 10.3p. 305Percent Composition and Chemical Formulas

  33. x 100 % = percent • All percent problems: part whole • Find mass of each element, • Divide by total mass of cmpd; & x 100 mass of element %mass of element = x 100% mass of cmpd

  34. % composition from mass • Calculate the percent composition of a compound that is made of 29.0 grams of Ag with 4.30 grams of S. 29.0 g Ag X 100 = 87.1 % Ag 33.3 g total Total = 100 % 4.30 g S X 100 = 12.9 % S 33.3 g total

  35. % comp from the chemical formula • If we know formula, assume you have 1 mole, • Subscripts used to calculate mass of each element in 1 mole of cmpd • sum of masses is molar mass Mass of element in 1 molcmpd x 100% %mass = Molar mass of cmpd

  36. % Composition Examples

  37. % composition as conversion factor • We can also use % as conversion factor to calculate # grams of element in cmpd • Calculate % C in C3H8 • What is mass of C in 82.0 g sample of propane (C3H8) 67.1 g C

  38. % Composition 4:15

  39. What is an Empirical Formula? • Like ingredients for recipe – double recipe, you double each ingredient, but ratio of ingredients stays same • Empirical formula: lowest whole number ratio of atoms in cmpd

  40. Calculating Empirical • Find lowest whole number ratio C6H12O6 CH4N • A formula is not just ratio ofatoms, it is also ratio ofmoles • 1moleculeof CO2 = 1atom of C and2 atoms of O • 1molof CO2 =1 mol Cand2 mol O = CH2O = this is already the lowest ratio.

  41. Calculating Empirical • get a ratio from % composition • Assume you have a 100 g sample - thepercentagebecomegrams (75.1% = 75.1 grams) • Convert gramstomoles. • Find lowest whole number ratio by dividing each # ofmolesby smallest value

  42. Example • Calculate empirical formula of cmpd composed of 38.67 % C, 16.22 % H, and 45.11 %N. • Assume 100 g sample, so • 38.67 g C x 1mol C = 3.22 mole C 12.0 g C • 16.22 g H x 1mol H = 16.22 mole H 1.0 g H • 45.11 g N x 1mol N = 3.22 mole N 14.0 g N CH5N Now divide each value by the smallest value

  43. Example • The ratio is 3.22 mol C = 1 mol C 3.22 mol N 1 mol N • The ratio is 16.22 mol H = 5 mol H 3.22 mol N 1 mol N = C1H5N1 which is = CH5N

  44. Practice Problem 36 p. 310

  45. What is a Molecular Formula? • Molecular formula: true # of atoms of each element in formula of cmpd • molecular cmpds only • Example: molecular formula for benzene is C6H6 (note that everything is divisible by 6) • Therefore, empirical formula =CH(the lowest whole number ratio)

  46. Formulas(continued) ionic compoundsALWAYS empirical (cannot be reduced). Examples: NaCl MgCl2 Al2(SO4)3 K2CO3

  47. 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)

  48. Empirical to molecular • Since empirical formula islowest ratio, the actual molecule weighs more Molar mass whole # to increase each coefficient in empirical formula = Empirical formula mass

  49. Empirical to molecular practice problem

  50. Empirical and Molecular Formulas 3:29

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