1. 1 Chapter 12 Chemical Quantities

2. 2 How do you measure things? How do you measure things? n We measure mass in grams. n We measure volume in liters. n We count atoms or compounds in MOLES.

3. 3 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. 6.02 x 10 23 is called Avogadro’s number. 6.02 x 10 23 is called Avogadro’s number. MEMORIZE this number!

4. 4 Types of questions n How many molecules of CO 2 are 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 ?

5. 5 Measuring Moles n The decimal number on the periodic table is also the mass of 1 mole of those atoms in grams. n Called molar mass n # on PT = 1 mole

6. 6 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?

7. 7 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

8. 8 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 molar mass of CH 4 is 16.05g What About Compounds?

9. 9 Molar Mass n The mass of one mole of a compound. n What is the molar mass of Fe 2 O 3 ? n 2 moles of Fe x 55.85 g = 111.70 g n 3 moles of O x 16.00 g = 48.00 g n The molar mass = 111.70 g + 48.00 g = 159.70g

10. 10 Examples Calculate the molar mass of the following: 1. Na 2 S 2. N 2 O 4 3. C 60 4. Ca(NO 3 ) 2 5. C 6 H 12 O 6 6. (NH 4 ) 3 PO 4

11. 11 Using Molar Mass Finding moles of compounds

12. 12 Molar Mass n The number of grams of 1 mole of atoms, ions, or molecules. n Make conversion factors n Change grams of a compound to moles of a compound.

13. 13 For example n How many moles is 5.69 g of NaOH?

14. 14 For example n How many moles is 5.69 g of NaOH?

15. 15 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles

16. 16 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l for NaOH:

17. 17 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l for NaOH: l 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g

18. 18 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l for NaOH: l 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g l 1 mole NaOH = 40.00 g

19. 19 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l for NaOH: l 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g l 1 mole NaOH = 40.00 g

20. 20 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l Need molar mass for NaOH l 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g l 1 mole NaOH = 40.00 g

21. 21 Examples 1. How many moles is 4.56 g of CO 2 ? 2. How many grams is 9.87 moles of H 2 O? 3. How many molecules in 6.8 g of CH 4 ? 4. 49 molecules of C 6 H 12 O 6 weighs how much?

22. 22 Gases and the Mole

23. 23 Gases n Many of the chemicals we deal with are gases. n Difficult to weigh n How do we know how many moles of gas we have?

24. 24 Gases Two things effect the volume of a gas: 1. Temperature 2. Pressure Compare at the same temperature and pressure.

25. 25 Standard Temperature and Pressure n STP is 0ºC and 1 atm pressure n At STP 1 mole of gas occupies 22.4 L n This is called molar volume

26. 26 Avogadro’s Hypothesis at the same temperature and pressure, equal volumes of gas have the same number of particles.

27. 27 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?

28. 28 Density of a Gas n The units will be g / L n We can determine the density of any gas at STP if we know its formula. n If you assume you have 1 mole, then the mass is the molar mass (P.T.) n At STP the volume is 22.4 L.

29. 29 Examples 1. Find the density of CO 2 at STP. 2. Find the density of CH 4 at STP.

30. 30 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 is the mass of 1 mole, since you have 22.4 L of a gas. n What is the molar mass of a gas with a density of 1.964 g/L? n 2.86 g/L?

31. 31 All the things we can change

32. 32 We have learned how to change n moles to grams n moles to atoms n moles to compounds n moles to liters n compounds to atoms n compounds to ions

33. 33 Moles Mass

34. 34 Moles Mass PT

35. 35 Moles Mass Volume PT

36. 36 Moles Mass Volume PT 22.4 L

37. 37 Moles Mass Volume Compounds PT 22.4 L

38. 38 6.02 x 10 23 Moles Mass Volume Compounds PT 22.4 L

39. 39 Moles Mass Volume Compounds 6.02 x 10 23 PT Atoms 22.4 L

40. 40 Moles Mass Volume Compounds 6.02 x 10 23 PT Atoms Ions 22.4 L

41. 41 Percent Composition Part x 100 % Part x 100 % Whole Whole n Find the mass of each component, then divide by the total mass.

42. 42 Example n Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

43. 43 Getting percent composition from the formula If we know the formula, assume you have 1 mole.

44. 44 Examples n Calculate the percent composition of C 2 H 4 n Calculate the percent composition of Aluminum carbonate.

45. 45 Empirical Formula From percentage to formula

46. 46 Empirical Formula - The lowest whole number ratio of elements in a compound. Molecular Formula - the actual ratio of elements in a compound.

47. 47 Empirical and Molecular Formula - 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

48. 48 Calculating Empirical Find the lowest whole number ratio n C 6 H 12 O 6 n CH 4 O 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.

49. 49 Calculating Empirical n We can get ratios from percent composition by assuming you have 100g. n The percentages become grams. n Can turn grams to moles. n Find lowest whole number ratio by dividing by the smallest.

50. 50 Example n Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N. n Assume 100 g so n 38.67 g C x 1mol C = 3.220 mole C 12.01 gC n 16.22 g H x 1mol H = 16.09 mole H 1.01 gH n 45.11 g N x 1mol N = 3.219 mole N 14.01 gN

51. 51 Example n The ratio is… 3.220 mol C = 1 mol C 3.219 molN 1 mol N 3.220 mol C = 1 mol C 3.219 molN 1 mol N n The ratio is… 16.09 mol H = 5 mol H 3.219 molN 1 mol N 16.09 mol H = 5 mol H 3.219 molN 1 mol N nC1H5N1nC1H5N1nC1H5N1nC1H5N1

52. 52 Example n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

53. 53 Empirical to molecular n empirical formula =lowest ratio n Actual molecule would weigh more by a whole number multiple. n Divide the actual molar mass by the mass of one mole of the empirical formula. n Caffeine has a molar mass of 194 g. what is its molecular formula?