2. 1 Chapter 8 Chemical Composition

3. 2 How you measure how much? How you measure how much? n You can measure mass, n or volume, n or you can count pieces. n We measure mass in grams. n We measure volume in liters. n We count pieces in MOLES.

4. 3 A. What is the Mole? n A counting number (like a dozen) n Avogadro’s number (N A ) n 1 mol = 6.02  10 23 items A large amount!!!!

5. 4 n 1 mole of hockey pucks would equal the mass of the moon! What is the Mole? n 1 mole of pennies would cover the Earth 1/4 mile deep! n 1 mole of basketballs would fill a bag the size of the earth!

6. 5 Moles n Defined as the number of carbon atoms in exactly 12 grams of carbon- 12. n 1 mole is 6.02 x 10 23 particles. n Treat it like a very large dozen n 6.02 x 10 23 is called Avagadro’s number.

7. 6 Representative particles n The smallest pieces of a substance. n For a molecular compound it is a molecule. n For an ionic compound it is a formula unit. n For an element it is an atom.

8. 7 Types of questions n How many oxygen atoms in the following? –CaCO 3 –Al 2 (SO 4 ) 3 n How many ions in the following? –CaCl 2 –NaOH –Al 2 (SO 4 ) 3

9. 8 Molar Mass n Mass of 1 mole of an element or compound. n Atomic mass tells the... – atomic mass units per atom (amu) – grams per mole (g/mol) n Round to 2 decimal places

10. 9 Molar Mass Examples n carbon n aluminum n zinc 12.01 g/mol 26.98 g/mol 65.39 g/mol

11. 10 Molar Mass Examples Molar Mass Examples n water n sodium chloride –H2O–H2O–H2O–H2O –2(1.01) + 16.00 = 18.02 g/mol –NaCl –22.99 + 35.45 = 58.44 g/mol

12. 11 Molar Mass Examples n sodium bicarbonate n sucrose –NaHCO 3 –22.99 + 1.01 + 12.01 + 3(16.00) = 84.01 g/mol –C 12 H 22 O 11 –12(12.01) + 22(1.01) + 11(16.00) = 342.34 g/mol

13. 12 Molar Conversions molar mass (g/mol) MASS IN GRAMS MOLES NUMBER OF PARTICLES 6.02  10 23 (particles/mol)

14. 13 Molar Conversion Examples Molar Conversion Examples n How many moles of carbon are in 26 g of carbon? 26 g C 1 mol C 12.01 g C = 2.2 mol C

15. 14 Molar Conversion Examples n How many molecules are in 2.50 moles of C 12 H 22 O 11 ? 2.50 mol 6.02  10 23 molecules 1 mol = 1.51  10 24 molecules molecules C 12 H 22 O 11 C 12 H 22 O 11

16. 15 Molar Conversion Examples Molar Conversion Examples n Find the mass of 2.1  10 24 molecules of NaHCO 3. 2.1  10 24 molecules 1 mol 1 mol 6.02  10 23 molecules = 290 g NaHCO 3 84.01 g 1 mol

17. 16 Types of questions n How many molecules of CO 2 are the in 4.56 moles of CO 2 ? n How many moles of water is 5.87 x 10 22 molecules? n How many atoms of carbon are there in 1.23 moles of C 6 H 12 O 6 ? n How many moles is 7.78 x 10 24 formula units of MgCl 2 ?

18. 17 Examples n How much would 2.34 moles of carbon weigh? n How many moles of magnesium in 24.31 g of Mg? n How many atoms of lithium in 1.00 g of Li? n How much would 3.45 x 10 22 atoms of U weigh?

19. 18 What about compounds? n in 1 mole of H 2 O molecules there are two moles of H atoms and 1 mole of O atoms n To find the mass of one mole of a compound –determine the moles of the elements they have –Find out how much they would weigh –add them up

20. 19 What about compounds? n What is the mass of one mole of CH 4 ? n 1 mole of C = 12.01 g n 4 mole of H x 1.01 g = 4.04g n 1 mole CH 4 = 12.01 + 4.04 = 16.05g n The Gram Molecular mass of CH 4 is 16.05g n The mass of one mole of a molecular compound.

21. 20 Gram Formula Mass n The mass of one mole of an ionic compound. n Calculated the same way.

22. 21 Molar Mass n The generic term for the mass of one mole. n The same as gram molecular mass, gram formula mass, and gram atomic mass.

23. 22 Examples n Calculate the molar mass of the following and tell me what type it is. n Na 2 S nN2O4nN2O4nN2O4nN2O4 nCnCnCnC n Ca(NO 3 ) 2 n C 6 H 12 O 6 n (NH 4 ) 3 PO 4

24. 23 Using Molar Mass Finding moles of compounds Counting pieces by weighing

25. 24 Molar Mass n The number of grams of 1 mole of atoms, ions, or molecules. n We can make conversion factors from these. n To change grams of a compound to moles of a compound.

26. 25 Examples n How many moles is 4.56 g of CO 2 ? n How many grams is 9.87 moles of H 2 O? n How many molecules in 6.8 g of CH 4 ? n 49 molecules of C 6 H 12 O 6 weighs how much?

27. 26 Gases and the Mole

28. 27 Gases n Many of the chemicals we deal with are gases. n They are difficult to weigh. n Need to know how many moles of gas we have. n Two things effect the volume of a gas n Temperature and pressure n Compare at the same temp. and pressure.

29. 28 Standard Temperature and Pressure n 0ºC and 1 atm pressure n abbreviated STP n At STP 1 mole of gas occupies 22.4 L n Called the molar volume n Avagadro’s Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles.

30. 29 Examples n What is the volume of 4.59 mole of CO 2 gas at STP? n How many moles is 5.67 L of O 2 at STP? n What is the volume of 8.8g of CH 4 gas at STP?

31. 30 Density of a gas n D = m /V n for a gas the units will be g / L n We can determine the density of any gas at STP if we know its formula. n To find the density we need the mass and the volume. n If you assume you have 1 mole than the mass is the molar mass (PT) n At STP the volume is 22.4 L.

32. 31 Examples n Find the density of CO 2 at STP. n Find the density of CH 4 at STP.

33. 32 The other way n Given the density, we can find the molar mass of the gas. n Again, pretend you have a mole at STP, so V = 22.4 L. n m = D x V n m is the mass of 1 mole, since you have 22.4 L of the stuff. n What is the molar mass of a gas with a density of 1.964 g/L? n 2.86 g/L?

34. 33 All the things we can change

35. 34 We have learned how to n change moles to grams n moles to atoms n moles to formula units n moles to molecules n moles to liters n molecules to atoms n formula units to atoms n formula units to ions

36. 35 Moles Mass

37. 36 Moles Mass PT

38. 37 Moles Mass Volume PT

39. 38 Moles Mass Volume PT 22.4 L

40. 39 Moles Mass Volume Representative Particles PT 22.4 L

41. 40 6.02 x 10 23 Moles Mass Volume Representative Particles PT 22.4 L

42. 41 Moles Mass Volume Representative Particles 6.02 x 10 23 PT Atoms 22.4 L

43. 42 Moles Mass Volume Representative Particles 6.02 x 10 23 PT Atoms Ions 22.4 L

44. 43 Percentage Composition Percentage Composition n the percentage by mass of each element in a compound

45. 44  100 = Percentage Composition %Cu = 127.10 g Cu 159.17 g Cu 2 S  100 = %S = 32.07 g S 159.17 g Cu 2 S 79.852% Cu 20.15% S n Find the % composition of Cu 2 S.

46. 45 %Fe = 28 g 36 g  100 = 78% Fe %O = 8.0 g 36 g  100 = 22% O n Find the percentage composition of a sample that is 28 g Fe and 8.0 g O. Percentage Composition

47. 46 n How many grams of copper are in a 38.0-gram sample of Cu 2 S? (38.0 g Cu 2 S)(0.79852) = 30.3 g Cu Cu 2 S is 79.852% Cu Percentage Composition

48. 47  100 = %H 2 O = 36.04 g 147.02 g 24.51% H 2 O n Find the mass percentage of water in calcium chloride dihydrate, CaCl 2 2H 2 O? Percentage Composition

49. 48 The Empirical Formula n The lowest whole number ratio of elements in a compound. n The molecular formula the actual ration of elements in a compound. n The two can be the same. n CH 2 empirical formula n C 2 H 4 molecular formula n C 3 H 6 molecular formula n H 2 O both

50. 49 Calculating Empirical n Just find the lowest whole number ratio n C 6 H 12 O 6 n CH 4 N n It is not just the ratio of atoms, it is also the ratio of moles of atoms. n In 1 mole of CO 2 there is 1 mole of carbon and 2 moles of oxygen. n In one molecule of CO 2 there is 1 atom of C and 2 atoms of O.

51. 50 Empirical Formula C2H6C2H6C2H6C2H6 CH 3 reduce subscripts n Smallest whole number ratio of atoms in a compound

52. 51 Calculating Empirical 1. Means we can get ratio from percent composition. 2. Assume you have a 100 g. The percentages become grams. 3. Find moles of each element. 4. Divide moles by the smallest # to find subscripts. 5. When necessary, multiply subscripts by 2, 3, or 4 to get whole #’s.

53. 52 Empirical Formula Empirical Formula n Find the empirical formula for a sample of 25.9% N and 74.1% O. 25.9 g 1 mol 14.01 g = 1.85 mol N 74.1 g 1 mol 16.00 g = 4.63 mol O 1.85 mol = 1 N = 2.5 O

54. 53 Empirical Formula N 1 O 2.5 Need to make the subscripts whole numbers  multiply by 2 N2O5N2O5N2O5N2O5

55. 54 Molecular Formula Molecular Formula n “True Formula” - the actual number of atoms in a compound CH 3 C2H6C2H6C2H6C2H6 empiricalformula molecularformula ?

56. 55 Molecular Formula 1. Find the empirical formula. 2. Find the empirical formula mass. 3. Divide the molecular mass by the empirical mass. 4. Multiply each subscript by the answer from step 3.

57. 56 Molecular Formula n The empirical formula for ethylene is CH 2. Find the molecular formula if the molecular mass is 28.1 g/mol? 28.1 g/mol 14.03 g/mol = 2.00 empirical mass = 14.03 g/mol (CH 2 ) 2  C 2 H 4

58. 57 Example n Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

59. 58 Examples n Calculate the percent composittion of C 2 H 4 ? n Aluminum carbonate.

60. 59 Example n Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N. n A compound is 43.64 % P and 56.36 % O. What is the empirical formula? n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula? n A compound is known to be composed of 71.65 % Cl, 24.27% C and 4.07% H. Its molar mas is known (from gas density) is known to be 98.96 g. What is its molecular formula?