1. 1 Chemical Quantities or

2. 2 How can you measure how much? How can 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.

3. 3 Examples n How much would 2.34 moles of carbon weigh?

4. 4 What is a mole equal to?

5. 5 Examples n How many moles of magnesium in 4.61 g of Mg?

6. 6 What about compounds? n 1 mole of H 2 O molecules is equal to the mass of 2 moles of H 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 REMEMBER: to go from moles-to- mass, you “mole-ti-ply”

7. 7 What about compounds? n What is the mass of one mole of CH 4 ?

8. 8 Molar Mass n What is the molar mass of Fe 2 O 3 ?

9. 9 Calculate the molar mass of the following n C 6 H 12 O 6

10. 10 n ammonium phosphate Molar Mass

11. 11 Using Molar Mass Finding moles of compounds Counting pieces by weighing

12. 12 Molar Mass n The number of grams in 1 mole of atoms, formula units, or molecules. n We can make conversion factors from these. n To change grams of a compound to moles of a compound. n Or moles to grams

13. 13 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

14. 14 Examples n How many moles in 72.1g C 2 H 4 ? n How many grams in 4.31 moles C 2 H 4 ?

15. 15 n How much (many?) is a mole? How much (many?) is a mole? How much (many?) is a mole?

16. 16 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 Avogadro's number.

17. 17 Representative particles n The smallest pieces of a substance. n For an element it is an atom. –Unless it is diatomic n For a molecular compound it is a molecule. n For an ionic compound it is a formula unit.

18. 18 Calculation question n How many moles of water is 5.87 x 10 22 molecules?

19. 19 Calculation question n How many molecules of CO 2 are the in 4.56 moles of CO 2 ?

20. 20 Calculation question n How many atoms of carbon are there in 1.23 moles of C 6 H 12 O 6 ?

21. 21 Examples n How many atoms of lithium in 1.00 g of Li?

22. 22 Examples n How much would 3.45 x 10 22 atoms of U weigh?

23. 23 Percent Composition n Like all percents n Part x 100 % whole n Find the mass of each component, n divide by the total mass.

24. 24 Getting it from the formula n If we know the formula, assume you have 1 mole. n Then you know the pieces and the whole.

25. 25 Examples n Calculate the percent composition of C 2 H 4 ? n What is the percent composition of aluminum carbonate.

26. 26 Example n Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

27. 27 Percent to Mass n Multiply % by the total mass to find the mass of that component. n How much aluminum in 450 g of aluminum carbonate?

28. 28 Empirical Formula From percentage to formula

29. 29 The Empirical Formula n The lowest whole number ratio of elements in a compound. n The molecular formula is the actual ratio of elements in a compound. n The two can be the same. n CH 2 is an empirical formula n C 2 H 4 is a molecular formula n C 3 H 6 is a molecular formula

30. 30 Finding Empirical Formulas n Just find the lowest whole number ratio n C 6 H 12 O 6 nC6H4N2nC6H4N2nC6H4N2nC6H4N2

31. 31 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 g/mol n 16.22 g H x 1mol H = 16.1 mole H 1.01 g/mol n 45.11 g N x 1mol N = 3.220 mole N 14.01 g/mol

32. 32 Example n Divide each of the derived values by the lowest value n 3.220 mol C = 1 3.220 mol C n 16.1 mol H = 5 3.220 mol C n 3.220 mol N = 1 3.220 mol C n Empirical formula is :C 1 H 5 N 1

33. 33 n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

34. 34 Empirical to molecular n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula? n Since the empirical formula is the lowest ratio the actual molecule would weigh the same or more. n By a whole number multiple. n Divide the actual molar mass by the the mass of one mole of the empirical formula. n You will get a whole number. n Multiply the empirical formula by this.

35. 35 Example n n A compound has an empirical formula of CH 2 O and a molar mass of 180.0 g/mol. What is its molecular formula?

36. 36 Example n n A compound has an empirical formula of ClCH 2 and a molar mass of 98.96 g/mol. What is its molecular formula?

37. 37 Percent to molecular n Take the percent x the molar mass –This gives you mass in one mole of the compound n Change this to moles –You will get whole numbers –These are the subscripts n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula?

38. 38 Gases and the Mole

39. 39 Gases n Many of the chemicals we deal with are gases. n They are difficult to weigh, so we’ll measure volume n Need to know how many moles of gas we have. n Two things affect the volume of a gas n Temperature and pressure n Compare at the same temp. and pressure.

40. 40 Standard Temperature and Pressure n Avogadro's Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles. n 0ºC and 1 atmosphere pressure n Abbreviated atm n 273 K and 101.3 kPa n kPa is kiloPascal

41. 41 At Standard Temperature and Pressure n abbreviated STP n At STP 1 mole of gas occupies 22.4 L n Called the molar volume n Used for conversion factors n Moles to Liter and L to mol

42. 42 Examples n What is the volume of 4.59 mole of CO 2 gas at STP?

43. 43 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.

44. 44 Examples n Find the density of CO 2 at STP.

45. 45 Quizdom n Find the density of CH 4 at STP.

46. 46 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?

47. 47 All the things we can change

48. 48 Volume Ions Atoms Representative Particles Mass PT Moles 6.02 x 10 23 22.4 L Count

49. 49 Example n n Ibuprofen is 75.69 % C, 8.80 % H, 15.51 % O, and has a molar mass of about 207 g/mol. What is its molecular formula?

50. 50 Calculating Empirical Formulas n Means we can get ratio from percent composition. n Assume you have a 100 g. n The percentages become grams. n Turn grams to moles. n Find lowest whole number ratio by dividing everything by the smallest moles.