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 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.

4. 4 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 or moles to grams

5. 5 Examples n How much would 2.34 moles of carbon weigh?

6. 6 Examples n How many moles of magnesium in 4.61 g of Mg?

7. 7 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”

8. 8 What about compounds? n What is the mass of one mole of CH 4 ? n 1 mole of C = 12 g n 4 mole of H x 1 g = 4g n 1 mole CH 4 = 12.01 + 4.04 = 16.05g

9. 9 Molar Mass n The mass of 1 mole n What is the molar mass of Fe 2 O 3 ? n 2 moles of Fe x 56 g = 112 g n 3 moles of O x 16 g = 48 g n The GFM = 112 g + 48 g = 160g

10. 10 Examples n How many grams in 4.31 moles C 2 H 4 ? n How many moles in 72.1g C 2 H 4 ? n ammonium phosphate

11. 11 n How much (many?) is a mole? How much (many?) is a mole? How much (many?) is a mole?

12. 12 Volume Ions Atoms Representative Particles Mass PT Moles 6.02 x 10 23 22.4 L Count

13. 13 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.

14. 14 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.

15. 15 Calculation question 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 lithium in 1.00 g of Li?

16. 16 Calculation question n How many atoms of carbon are there in 1.23 moles of C 6 H 12 O 6 ? n How much would 3.45 x 10 22 atoms of U weigh?

17. 17 Examples n How much would 3.45 x 10 22 atoms of U weigh?

18. 18 Percent Composition n Like all percents n Part x 100 % whole n Find the mass of each component, n divide by the total mass.

19. 19 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.

20. 20 Examples n Calculate the percent composition of C 2 H 4 ? n What is the percent composition of aluminum carbonate.

21. 21 Example n Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

22. 22 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?

23. 23 Empirical Formula From percentage to formula

24. 24 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

25. 25 Finding Empirical Formulas n Just find the lowest whole number ratio n C 6 H 12 O 6 nC6H4N2nC6H4N2nC6H4N2nC6H4N2

26. 26 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

27. 27 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

28. 28 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.

29. 29 n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

30. 30 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?

31. 31 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?

32. 32 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?

33. 33 Gases and the Mole

34. 34 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.

35. 35 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

36. 36 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

37. 37 Examples n What is the volume of 4.59 mole of CO 2 gas at STP?