1. Chapter 10 The mole

2. 10.1 The Mole What do you ask for when you buy: 2 shoes 12 eggs 48 doughnuts 500 sheets of paper 1 pair 1 dozen 4 dozen 1 ream

3. Pair, Dozen, Ream These are all ways to batch a group of objects to make them easier to count! The object may change Eggs to Doughnuts But the number they represent is always constant!

4. What do eggs have to do with Chemistry? How many carbon atoms are in a teaspoon of carbon? 200,666,666,666,666,666,666,667 atoms!!! Can you even pronounce this?

5. 200,666,666,666,666,666,666,667 atoms!!! It would be nice if chemists had a batch like a dozen (but muuuucchh bigger) to handle this kind of number!

6. Meet the Mole! He is the Chemist’s “dozen”!

7. The chemist’s “dozen” is called the: MOLE (or the unit mol) 1 dozen atoms = 12 atoms 1 mole atoms = 602213670000000000000000 atoms How many moles are in 1 teaspoon of carbon atoms?.33 moles Or 1 mole = 6.02  10 23 What is a mole?

8. 11.0 Measuring Matter – Moles and Avogadro’s number Mole (mol) – the amount of a substance that contains the same number of particles as the number of atoms in 12 g of carbon-12. Like doughnuts are counted in dozens, the mole is a SI unit for counting the amount of a substance. –1 dozen pencils have the same number of particles as 1 dozen doughnuts –1 mole of carbon atoms have the same number of particles as 1 mole of water molecules This does not mean they both weigh the same, only that they have the same number of units or particles.

9. 10.1 Measuring Matter – Moles and Avogadro’s number Avogadro’s Number – the number of particles (6.02  10 23 ) in exactly one mole of a pure substance. –1 mole pencils = 6.02  10 23 pencils –1 mole water molecules = 6.02  10 23 molecules of water 1 mole of Guaca-”mole”

10. 10.1 Measuring Matter – Moles and Avogadro’s number If Avogadro’s number is the number of particles in 1 mole, how do you know what kind of particle you have? Remember… Particles can be an atom, formula unit, or molecule?

11. 10.1 What kind of particle? Particles can be an atom, molecule, or formula unit? Atom – one atom Formula unit - starts with a metal or NH 4 + Molecule – ALL nonmetal atoms

12. 10.1 Measuring Matter – Moles and Avogadro’s number How do you know what kind of particle you have? Examples 1.NaCl 2.H 2 O 3.H 2 4.Na formula unit (form.unit) molecule atom

13. 10.1 Measuring Matter – Moles and Avogadro’s number Avogadro’s number is really an equality! 1 mole C = 6.02  10 23 atoms C What can you use equalities to do? Equalities are conversion factors in Dimensional Analysis problems!

14. 10.1 Measuring Matter – Moles and Avogadro’s number How many atoms are in 3.2 mol of C? = atoms C mol C atoms C 1 6.02  10 23 1.9  10 24

15. 10.1 Measuring Matter – Moles and Avogadro’s number How many moles of water molecules are in 3.7  10 21 molecules of H 2 O? = mol H 2 O mol H 2 O Molecules H 2 O 1 6.02  10 23.0061 3.7  10 21 molecules H 2 O

16. Why did Avogadro stop going to the chiropractor on October 24 th ? He was only tense to the 23 rd ! (10 23 )

17. 10.1 Mass and the Mole Molar mass – the mass, in grams, of one mole of a substance. How do you find this mass? The atomic mass printed on the periodic table has two meanings! 1.It is the average mass of one atom in atomic mass units (amu or u). 2.It is also the mass of one mole of atoms of a given element in grams.

18. 10.1 Mass and the Mole Finding molar mass –Atoms – the mass of 1 mole of any atom is the same as the atomic mass in grams. (Round your masses to the tenths) 1mole H atoms = 1.008 g 1mole C =  1.0 g 12.011 g  12.0 g 1mol Cu = 63.546 g  63.5 g

19. 10.1 Mass and the Mole Molar mass continued –Compounds – the mass of 1 mole of a compound is the sum of the masses of the atoms. 1mole H 2 O = 2 H = 1 O = 2(1.0g) = 2.0g 16.0g 2.0g + 16.0g = 18.0g 18.0g As a molar mass, it would be expressed as 18.0 g/mol

20. 10.1 Mass and the Mole Molar mass is really an equality! For water: 1mole H 2 O = 18.0 g What can you use equalities to do? Equalities are conversion factors in Dimensional Analysis problems!

21. 10.1 Mass and the Mole How many moles are in 242 g of water? mol H 2 O g H 2 O 1 18.0 13.4 242 g H 2 O = mol H 2 O

22. 10.1 Mass and the Mole What is the mass of 3.77 mol of gold? g Au mol Au 197.0 1 743 3.77 mol Au = g Au

23. Conversion Review Moles (mol) Mass (g) Particles (atom/molecule/ formula unit) Molar mass Avogadro’s #

24. Moles (mol) Mass (g) Particles (atom/molecule/ formula unit) Molar mass Avogadro’s # 1.Write Given and Target with 2 part units. 2.Use the flow chart to see how many steps. 3.Cancel the unit you don’t want. NO NUMBERS YET!!!!! 4.Use the flow chart to see what unit to put on top. 5.Add numbers to make them equal. 6.Continue until done. mol H 2 O g H 2 O 1 18.0 8.09  10 24 242 g H 2 O = molecules H 2 O mol H 2 O molecule H 2 O 6.02  10 23 1 How many water molecules are in 242 g of water?

25. 10.3 Percent Composition How would you calculate the percent females in this room? All percents are calculated in the same way!

26. 10.3 Percent Composition Percent Composition – the percent by mass of each element in a compound. molar mass

27. What is the percent composition of each element in water? Mass H = Mass O = 16.0 g 10.3 Percent Composition 2(1.0 g) = 2.0 g + 18.0 g % H = 11% H% O = 88.9% O

28. Using Percent Composition If a glass of water contains 648 g of water, how many grams of hydrogen would it hold? Remember, water is 11% hydrogen What is 11% of 648 g? 71.28 g 71 g H

29. C H N O 81042 Where are the missing carbons?

30. 10.3 Empirical Formula What does H 2 O mean? Does it mean 2 atoms of H for every atom of O? YES Does it mean 2 g of H for every 1 g of O? NEVER

31. 10.3 Empirical Formula What does H 2 O mean? Does it mean 2 moles of H for every mole of O? Always! Formulas are not only ratios of atoms, they are also ratios of MOLES

32. Formulas are ratios of moles!

33. 10.3 Empirical Formula Empirical Formula – simplest whole number ratio of moles of the atoms in a substance. Experimental method that is the first step in finding the formula of a compound. Circle the empirical formulas! H 2 SO 4 H2OH2O N2O4N2O4 C2H6C2H6 NO 2 C 6 H 12 O 6 NaCl

34. 10.3 Empirical Formula Finding the empirical formula 1.Find the mass of each element in the compound. –Usually given –If given as %, then change % to g. 36% H and 64% C  36 g H and 64 g C 2.Convert masses to moles. –Use molar masses. (Don’t worry about significant digits in this step!) 36 g H = mol H g H mol H1 1.0 36

35. 10.3 Empirical Formula 3.Find the smallest whole number ratio of moles. a. Write the results of step 2 like a formula. If C = 1.2 mol and H = 4.8 mol C 1.2 H 4.8 b. Divide by the smallest mole amount. C 1.2 H 4.8 c. If not all whole numbers, multiply by 2,3, or 4 … 1.2 CH 4

36. 10.3 Empirical Formula Examples for step 3 X =.029 molX =.009 mol Y =.039 molY =.006 mol X.029 Y.039.029 X 1 Y 1.34 Multiply by 3 X3Y4X3Y4 X.009 Y.006.006.006 X 1.5 Y 1 Multiply by 2 X3Y2X3Y2

37. 10.3 Empirical Formula More Examples for step 3 X = 1.47 molX = 2.4  10 -4 mol Y = 3.68 molY = 7.3  10 -4 mol X 1.47 Y 3.68 1.47 X 1 Y 2.5 Multiply by 2 X2Y5X2Y5 2.4  10 -4 X 1 Y 3.04 X1Y3X1Y3

38. 10.3 Empirical Formula - Example A sample of an unknown gas contains 43.2 g of carbon and 115.8 g of oxygen. What is the empirical formula? 1.Find Masses 43.2 g C115.8 g O 2.Change to moles 43.2 g C = mol C g C mol C1 12.0 3.60 115.8 g O = mol O g O mol O1 16.0 7.24

39. 10.3 Empirical Formula - Example 3.Get whole numbers C 3.60 O 7.24 3.60 C 1 O 2.01 CO 2

40. “Find the mass, Convert to moles. Divide by small, Multiply till whole.”

41. 10.3 Molecular Formula Molecular formula - is some whole number multiple of the empirical formula. –HO is an empirical formula –H 2 O 2 is twice HO –(HO) X and X = 2 –For C 6 H 12 O 6, the empirical formula is CH 2 O and X=6 To convert an empirical formula to a molecular formula you must find X.

42. 10.3 Molecular Formula - Example An unknown gas is found to have an empirical formula of NO 2 and a molar mass of 92.0 g/mol. What is the molecular formula? Molecular formula = (NO 2 ) X = (NO 2 ) 2 = N 2 O 4

43. The Formula of a Hydrate Hydrate: a compound that has a specific number of water molecules bound to its atoms. –symbolized by “Compound. ? H 2 O” Examples: CuSO 4. 5 H 2 O and BaCl 2. 2 H 2 O –Naming: Name compound, then use prefixes to describe number of water molecules followed by hydrate Example: sodium carbonate decahydrate Na 2 CO 3. 10 H 2 O deca- means 10, hydrate refers to water

44. The Formula of a Hydrate Anhydrous: “without water” –When a hydrate is heated, it drives off the water and the compound is no longer hydrated. This leaves an anhydrous substance. –Name and Formula of anhydrous compounds Same as the ionic compound –Anhydrous sodium carbonate is usually just called sodium carbonate –Na 2 CO 3

45. The Formula of a Hydrate Finding the Formula of a Hydrated Crystal The formula of a hydrated compound is really a type of empirical formula. It’s the smallest whole number ratio of water to 1 mole of the compound –Na 2 CO 3. 10 H 2 O –1 mole of Na 2 CO 3 for every 10 moles of water

46. The Formula of a Hydrate Finding the Formula of a Hydrated Crystal Like any empirical formula, you must: 1.Get mass Need to find mass of anhydrate and water 2.Get moles Convert masses → moles using molar masses 3.Get whole numbers Divide the moles of water by the moles of anhydrate (there is always the same or more water) and you will NEVER need to double or triple.

47. Hydrate Lab Lab: You will be analyzing a hydrated crystal to determine the formula of the hydrate. The substance will start as a hydrated crystal and will be measured in a crucible with a lid (always weigh it with the lid). –Weigh the empty crucible –Weigh the crucible with hydrate –The difference between these two would be the mass of the hydrated crystal, but we don’t need to calculate this!

48. Hydrate Lab Lab: After it is heated enough, all water will be driven off and you’ll be left with the anhydrate. –Weigh the crucible with anhydrate

49. Hydrate Lab Lab: When doing the lab, you will need to obtain all three masses! From these measurements, we can calculate the 2 masses we need to find the formula of the hydrated crystal – the mass of the anhydrate and the mass of the water. –Mass of anhydrate Mass crucible w/anhydrate – mass crucible –Mass of water Mass crucible w/hydrate – mass crucible w/anhydrate

50. Lab Data Mass of crucible & lid: 11.36g Mass of crucible, lid, & hydrated copper sulfate: 13.86g Mass of crucible, lid, & anhydrous copper sulfate: 12.95g

51. 1A. Calculate mass (grams) of anhydrous copper sulfate. (mass crucible,lid, anhydrous copper sulfate) - (mass crucible & lid) 12.95 g – 11.36 g = 1.59 g 1B. Calculate mass of water. (mass crucible,lid, hydrated copper sulfate – mass crucible,lid, anhydrous copper sulfate) 13.86 g – 12.95 g = 0.91 g

52. 2A. Calculate moles anhydrous CuSO 4 –Grams → moles using molar mass 1.59g CuSO 4 159.6g CuSO 4 1 mol CuSO 4 = 0.00996 mol CuSO 4 Lab Calculations Mass of anhydrate: 1.59 g Mass of water: 0.91 g Moles of anhydrate: Moles of water: 0.00996 mol CuSO 4 2B. Calculate moles of water. –Grams → moles using molar mass.91g H 2 O 18.0 g H 2 O 1 mol H 2 O = 0.050 mol H 2 O

53. Lab Calculations Mass of anhydrate: 1.59 g Mass of water: 0.91 g Moles of anhydrate: 0.00996 mol Moles of water: 0.050 mol

54. What is the formula of the hydrate? The ratio of moles of H 2 O to moles of CuSO 4, n, is the coefficient that precedes H 2 O in the formula for the hydrate. Formula : CuSO 4 ∙ __H 2 O 5