1. Chemical Quantities (the MOLE) Chapter 10

2. Counting Units  How many is a dozen?  How many does the word “couple” stand for?  How many sheets are in a ream of paper?  Answers: 12, 2, 500

3. Chemist’s “Dozen”  Chemists also have a counting unit for a certain number of things  A mole = Avogadro’s number of things  1 mole = 6.02 x 10 23 things –Could be atoms, ions, particles, molecules, etc.  How big is that number? –See Sheet: A mole is a lot of things

4. Sample Problem  How many atoms are contained in 4.5 moles of copper?  How many moles are in 8.3 x 10 24 molecules of water?  Answers: 2.71 x 10 24 atoms, 13.8 moles

5. Masses and Moles  If you had a dozen bowling balls and a dozen ping pong balls, which would have more mass? Why?  The bowling balls would have more mass because each individual bowling ball has more mass than a ping pong ball.

6. Mass of Elements and Moles  Which would be heavier: a mole of copper or a mole of lead? Why?  A mole of lead would be heavier because an atom of lead is heavier than an atom of copper (64 amu vs. 207 amu)

7. Carbon-12 and Avogadro’s Number  Avogadro’s number is related to the mass of an element, specifically carbon (the isotope carbon-12)  Used exactly 12 grams of carbon and determined the number of atoms in that sample  There are 6.02 x 10 23 atoms in 12 grams of carbon (1 mole)

8. Definition of Mole  A mole is the amount of matter that contains as many objects (atoms, molecules, etc.) as the number of atoms in exactly 12 g of 12 C. –Mole connects amu & gram = “Chemists dozen”

9. Molar Mass  How do you determine moles if you have a compound?  What is the mass of 1 mole of NaCl?  Answer: 58 g/mol  Molar Mass: the mass of one mole of a substance (element or compound)

10. Sample Problem- molar mass  What is the mass of one mole of –water? –Carbon dioxide? –C 6 H 12 O 6 (glucose) –Ca(NO 3 ) 2 –Answers: 18 g/ mol, 44, 180, 164

11. Sample Problem- moles to mass  How many grams of sodium chloride (NaCl) did you use if you put 0.250 moles of it on your French fries?  Answer: 14.5 g

12. Sample Problem – mass to moles  How many moles of helium are in a balloon that contains 4.5 g of helium?  Answer: 1.1 moles

13. Molar Volume  Can you determine moles from the volume of a substance? –For solids and liquids: must use density –For gases: related to the temperature and pressure of the gas.

14. Avogadro’s Hypothesis  Equal volumes of gases at the same temperature and pressure contain the same number of molecules  All gases behave the same way, no matter the chemical identity

15. Standard Temp and Pressure  Because the volume of a gas depends on the temp and pressure of the gas, a standard reference is necessary  STP: Standard Temperature and Pressure –Std Temp: 0 o C or 273 K –Std Press: 101.3 kPa or 1 atm  1 mole of a gas at STP = 22.4 L

16. Sample Problem – Volume to moles  How many moles of helium are in a balloon has a volume of 6.40 L? (at STP)  Answer:.286 moles

17. Study Buddy Review  What is Avogadro’s Number?  What is a mole?  What is molar mass?  How do you find the molar mass of a compound?  What is STP?

18. Percent Composition

19. Penny Composition  Remember: a penny’s composition has changed through the years –Most recent change 1982  If a post-1982 penny is 2.500 g and it has 0.0625 g copper, what percentage is copper? What percentage is zinc?  Answers: 2.50 % Cu, 97.50 % Zn

20. Percent Composition  mass of each element in compound compared to the entire mass of the compound, multiplied by 100%

21. Sample Problem  What is the percentage composition of copper in chalcocite, Cu 2 S?  Answer: 80 %

22. Consider these  C 6 H 12 O 6 glucose (blood), fructose (corn syrup)  C 5 H 10 O 5 ribose (RNA), xylose  What do the formulas have in common?

23. Empirical formulas  CH 2 Osmallest whole number ratio  Carbohydrates: C n (H 2 O) n  Empirical Formulas: smallest whole number ratio of the atoms in a substance

24. Molecular Formulas  Molecular Formula: whole number multiple of empirical – –tells the actual number of atoms in a substance as it exists in nature –Can be the same as the empirical formula