1. The Mole Honors Chem

2. -How do we measure chemical quantities? -What units of measure do we use?

3. Atomic Mass Unit (amu): Relative mass based on 12 C atom 1 amu = 1/12 the mass of a 12 C atom

4. An amu is very small 1 amu is about 1 x 10 -24 grams Not practical

5. How can grams be related to amu’s?

6. # atoms H O Ratio 11amu 16amu 1:16 22amu 32amu “ 1 dozen 12amu 192amu “ A gross144amu 2304amu “ 6.02x10 23 6.02x10 23 9.63x10 24 “

7. Relating amu’s to grams 6.02x10 23 atoms of any element will have a mass equal to the element’s atomic mass in grams

8. For example 6.02x10 23 atoms C = 12 g 6.02x10 23 atoms Na = 23 g 6.02x10 23 atoms Cu = 63.5 g

9. This number 6.02x10 23 is called Avogadro’s number. THE MOLE!

10. THE MOLE Is a number: –6.02x10 23 of anything is a mole Is a mass: –The mass of an element expressed in grams –The mass of a compound expressed in grams

11. Examples 12 g is a mole of carbon atoms, C How many atoms are in this mass? 32 g is a mole oxygen molecules, O 2 How many molecules are in this mass?

12. Some terms

13. Atomic Mass: The weighted average of the masses of all the isotopes of an element.

14. Molar Mass: The mass in grams of one mole of a substance. Other names used for molar mass are molecular weight gram formula mass

15. Problems 1.What is the mass of one mole of carbon dioxide? 44 g 2. Find the number of moles in 185 g Calcium hydroxide? 2.50 mol 3. How many grams are in 0.450 moles of water? 8.10 g

16. 4. How many copper atoms are in 3.00 moles of copper? 1.8x10 24 atoms 5. What is the mass of 3.01x10 24 molecules of water? 90.0 g

17. 6. How many molecules are in a 20.0 g sample of carbon dioxide? 2.74x10 23 molecules

18. 7. For a sample of 150.0 g of H 2 SO 4, calculate: a)The number of moles of formula units of H 2 SO 4 present. b) The number of oxygen atoms present. 1.53 mol, 3.68 x10 24 atoms O

19. Moles in Solution  Molarity  Expresses concentration  Ratio of particles of solute per Liter(dm 3 ) of solution  Moles/L or moles/dm 3 MM

20. Two types of calculations Example 1: finding Molarity What is the molarity of a 250. cm 3 solution containing 3.03 g potassium nitrate? 0.120M

21. Example 2: Finding amount of solute needed to prepare a solution. How many grams of glucose must be dissolved in 325 mL to make a 0.258 M solution? 15.1 g glucose

22. Empirical formula Atoms combine in whole number ratios Moles of atoms will also be in whole number ratios The mole ratios can be used to calculate an empirical formula

23. Examples 1. A compound was found to have a 2:1 mole ratio of oxygen to carbon. What is the formula? CO 2

24. 2. An analysis showed a compound sample consists of 64.0 g oxygen and 24.0 g carbon. What is the formula? CO 2

25. 3. What is the empirical formula of a compound containing 65.2% arsenic and 34.8% oxygen by mass? As 2 O 5

26. A rhyme to help remember % to mass Mass to mole Divide by smallest (Multiply for whole)

27. In the text book Page 212 #59 #58

28. Molecular Formulas Contain the actual number of atoms of each element Is a whole number multiple of the empirical formula The molecular mass must be known

29. Examples 1. A compound has an empirical formula CH. What is the molecular formula if the molar mass is 78.0 g/mol? C 6 H 6

30. 2. What is the molecular formula of a compound that has a molar mass of 30.0 amu if analysis shows 20.0% hydrogen and 80.0% carbon by mass? hint: find empirical formula first. C2H6C2H6

31. 3. Find the empirical and molecular formula for a compound containing 2.80 g nitrogen and 6.40 g oxygen? The molecular mass is 46.0 g/mol NO 2 is both empirical and molecular

32. Hydrates Salts with water chemically bound to them CoCl 2. 6H 2 O I mole of CoCl 2 with 6 moles of H 2 O

33. Calculation The mass of the hydrate of CuSO 4 is 2.00g. After heating,.710 g of water were lost. What is the formula of the hydrate?