How do we use the mole?

1. How do we use the mole? Chemistry Unit 9

2. Main Ideas • Chemists use the mole to count atoms, molecules, ions and formula units. • A mole always contains the same number of particles, however, moles of different substances have different masses. • The molar mass of a compound can be calculated from its chemical formula and can be used to convert from mass to moles of that compound. • A molecular formula of a compound is a whole-number multiple of its empirical formula

3. Measuring matter

4. 9.1 Measuring matter: Objectives • Explain how a mole is used to indirectly count the number of particles of matter. • Relate the mole to a common everyday counting unit. • Convert between moles and number of representative particles.

5. Mole The mole is the SI base unit for measure of amount of a substance: 6.0221367 x 1023 • The number of carbon atoms in exactly 12 g of pure carbon-12. • Called Avogadro’s number – Italian physicist who in 1811, determined the volume of 1 mol of gas. • By mass, we can determine the number of particles (atoms, molecules) in a sample. • We typically round to 3 sig figs – 6.02 x 1023

6. The mole: A good comparison The mole is a number. What other unit is used in a similar manner? • A dozen flowers, doughnuts or eggs. • A baker’s dozen of cookies or bagels. • A pair of socks or friends If you have a dozen flowers and a dozen eggs, do they weigh the same?

7. The mole as a Conversion factor In order to convert between moles and number of particles we need to use this ratio of equivalent values (conversion factor) to express the same quantity in different units.

8. Example Problems 1&2 How many particles are in 3.5 mols? How many moles of atoms are in 9.63 x 1026atoms?

9. Question? What does the mole measure? A.mass of a substance B.amount of a substance C.volume of a gas D.density of a gas

10. Question? What is the conversion factor for determining the number of moles of a substance from a known number of particles? A. B. C.1 particle 6.02  1023 D.1 mol  6.02  1023particles

11. Practice problems • Page 322 #1-4; page 324 #5-14

12. mass and the mole

13. Mass and the Mole: objectives • Relate the mass of an atom to the mass of a mole of atoms. • Convert between number of moles and the mass of an element. • Convert between number of moles and number of atoms of an element.

14. Molar mass Molar mass is the mass in grams of one mole of any pure substance. • Units are given in g/mol • Mass of the periodic table is given in amu, but also g/mol

15. Example Problems 3&4 If I need 3 mols of Cu, how do I measure the amount? I measured 5.0g of Iron, how many atoms do I have?

16. Conversions

17. Example Problems 5&6 How many atoms of gold are in a U.S. Eaglebullion coin with a mass of 31.1g? How much does 5.8 x 1015 atoms of lead weigh?

18. question The mass in grams of 1 mol of any pure substance is: A.molar mass B.Avogadro’s number C.atomic mass D.1 g/mol

19. question Molar mass is used to convert what? A.mass to moles B.moles to mass C.atomic weight D.particles

20. Practice problems • Page 328 #15-16; page 329 #17-18 • Page 331 #19-21; page 332 #22-27

21. moles of compounds

22. Moles of Compounds: Objectives • Recognize the mole relationships shown by a chemical formula • Calculate the molar mass of a compound. • Convert between number of moles and mass of a compound. • Apply conversion factors to determine the number of atoms or ions in a known mass of a compound

23. Calculate molar mass Steps to calculate molar mass: • Count the number of atoms in each molecule. • Find the molar mass of each atom. • Multiply the molar mass of each atom to the number of atoms in a compound. • Add the total molar masses together.

24. Example Problem 7 Find the molar mass of the following compounds/molecules. • H2O • NaCl • H2SO4

25. Example Problem 7 cont.. • Al2O3 • Fe2(SO4)3 • CCl2F2

26. Number of Atoms To determine the number of atoms or ions in a known mass of a compound • Find the molar mass of the compound. • Use molar mass and the mole as conversion factors to get the units needed.

27. Example Problem 8 What is the mass of 2.5 mols of (C3H5)2 S?

28. example problem 9 Calculate the number of moles of Ca(OH)2 in 325g of the compound?

29. example problem 10 How many atoms are in 212g of water?

30. question How many moles of OH— ions are in 2.50 moles of Ca(OH)2? A.2.00 B.2.50 C.4.00 D.5.00

31. question How many particles of Mg are in 10 moles of MgBr2? A.6.02 x 1023 B.6.02 x 1024 C.1.20 x 1024 D.1.20 x 1025

32. Practice problems • Page 335 #29-36; page 336 #37-41; page 339 #42-46

33. empirical and molecular formulas

34. Empirical and Molecular formulas: Objectives • Explain what is meant by the percent composition of a compound. • Determine the empirical and molecular formulas for a compound from percent and actual mass data. • Explain what a hydrate is and relate the name of the hydrate to its composition. • Determine the formula of a hydrate from laboratory data.

35. Percent composition The percent composition is a percent by mass of each element in a compound. Steps to determine percent composition of a compound: • Assume 1 mole of a compound. • Calculate molar mass of each element in the compound. • Use each element’s molar mass to calculate percent by mass.

36. Percent by mass Percent by mass is a description of the amount of an element in a compound. • Percent by mass =

37. Example problem 11 What is the percent by mass of each element in NaHCO3?

38. Empirical formula The empirical formula is the smallest whole number ratio of elements in a compound • This ratio provides the subscripts for the elements. • May or may not be the same as the actual molecular formula. • If they are different the molecular formula will be a simple multiple of the empirical formula. • Hydrogen peroxide: HO- empirical formula H2O2 – actual formula (molecular formula)

39. Empirical formula Steps to figure empirical formula from percent composition: • Assume an overall 100g sample of the compound. • Each element’s percentage can be used as mass in calculations. • Use this ‘mass’ to convert to moles. This provides a ‘mole ratio’ for the compound.

40. Empirical formula • Since these mole ratios are not whole numbers, we convert them to whole numbers what can be used as subscripts by dividing them all by the smallest ratio. (We assume the smallest mole ratio is a 1 in the compound)

41. example Problem 12 A compound has the following mass percentages: C – 48.64%, H – 8.16%, O – 43.20% What is the empirical formula for this molecule?

42. Molecular formula The molecular formula specifies the actual number of atoms of each element in one molecule/formula unit of the substance.

43. Molecular formula Steps to determine the molecular formula: • Determine the molar mass of the empirical formula. • Determine the molar mass of the actual compound. (might be given to you) • Divide the molar mass of the actual compound by the molar mass of the empirical formula • Multiply all subscripts of the empirical formula by this molar mass ratio.

45. Example Problem 13 The mass of benzene has been experimentally determined to be 78.12 g. We know that benzene is 92 % C by mass and 8 % H by mass. What is the molecular formula of benzene?

46. Practice Problems • Page 344 #54-57 • Page 348 #58-61 • Page 350 #62-66

47. hydrates Hydrates are solid ionic compounds in which water molecules are trapped. • Hydrates are formed when water molecules adhere to the ions as the solid forms. • Water molecules become a part of the crystal solid structure. • The number of water molecules associated with each molecule is written following a dot after the molecular formula: • Na2CO310H2O

48. hydrates • Names of these compounds are named with a prefix representing the number of water molecules and the word hydrate. • Na2CO3  10H2O – sodium carbonate decahydrate • Prefixes are the same as the ones used in naming covalent molecules

49. naming hydrates

50. Anhydrous An anhydrous is a compound without water. • When a hydrate is heated, water molecules are driven off leaving the compound.