2. 1 mole  1 mole marbles = covers Earth to depth of 50 miles

3. Counting by weighing Purpose: Calculate the amount of pennies in the bag by weighing them. ( You may take the pennies out of the bag to weigh them, but do not count them) Data Calculations  Atoms are too small to count so we weigh an amount (mole) and calculate the number

4. Avogodro  The actual size of the molecule is not important  Avagadro concluded that equal volumes of gas have equal number of molecules By the way, nobody really knew what a molecule was at the time

5. Some words mean numbers  Pair  Dozen  Baker’s dozen  Gross  Ream

6. Mole 1 mole = 6.02 x 10 23 representative particles (atoms) (molecules) (formula units) ( the number of carbon-12 atoms in 12.00 g)

7. Atomic mass units  We measure atoms in atomic mass units  Atomic mass unit = AMU = 1.66x10 -27 Kg  It is defined as 12 for Carbon – 12  Carbon 12 has six protons and six neutrons 1 proton=1 amu (1.0078amu) 1 neutron= 1 amu (1.0087 amu)

8. Chemical Measurements  Atomic Mass The weighted average of all the mass numbers for all the isotopes of the atom (a.m.u.)  Formula Mass The sum of all the atomic masses for all atoms in the compound. (a.m.u.)

9. Calculate the atomic mass or formula masses  Na  Cl  Br 2  NaCl  CO 2  Mg(OH) 2 22.99 amu 35.45 amu 159.80 amu 58.45 amu 44.01 amu 58.33 amu

10. Molar Mass =The mass of 1 mole ( 6.02 x 10 23 ) of representative particles =The atomic mass in g =The formula mass in g

11. Atom, Molecule and Ions  Representative particles – smallest particle of that substance Substance Representative Particle element atom molcular compound molecule ionic compoundformula unit

12. Calculate the Molar mass  Ca H2H2  KNO 3  (NH 4 ) 2 S

13. Remember 1 mole = 6.02 x 10 23 r.p. 1 mole= molar mass(g)

14. practice 1. Determine the number of FU in.866 moles of AgNO 3 ? 2. Find the mass of 0.98 moles of CaCl 2.

15. Molar Volume The volume of 1mole of a gas at standard temperature – 0 o C standard pressure – 1 atmosphere (pressure at sea level) 1 mole gas at STP = 22.4 L

16. Remember 1 mole = 6.02 x 10 23 r.p. 1 mole= molar mass(g) 1 mole=22.4 L gas

17. Multistep problems  How many fu in 18.9 g NaCl?  How many L in 4.5 x 10 13 molecules Ne?

18. Solution= solute + solvent %mass= g solute g of solution PPM = parts of solute 1,000,000 parts of solution Molarity = Moles solute  L of solution

19. Calculate  1.What is the molarity if 13.5 g NaCl in 451 ml solution?  2.How many g of KOH are in 3.5 L of a.67 M solution?  3.How many ml of solution would you need to have 17.5 g of NaOH in a.35 M solution?

20. Percent Composition % element = g element x 100% g total compound

21. Formulas  Empirical simplest whole number ratio  Molecular Formula actual number of atoms in the formula

22. Practice: A compound has 13.5 g Ca 10.8 g O.675 g H What is its empirical formula?

23. Try this one  What is the empirical formula that is  25.9% N, 74.1% O

24. Spark  1.How many moles of HCl are there in 1.00 liters of 2.0 molar HCl?  2.How many milliliters of 13.2 g of HCL of a 3.4 M solution?  3.How many oxygen atoms are in 150.2 mL of 2.00 M H 2 SO 4 ?  4.How many sulfur atoms are in 158.2 g of aluminum thiosulfate, Al 2 (S 2 O 3 ) 3 ?

26. Molecular formula Set up table: EmpircalMolecuar formula Molar mass

27. A Chemical Reaction  Balance the following reaction  Hydrogen and oxygen react to produce water H 2 + O 2  H 2 O 2H 2 + O 2  2H 2 O

28. Why Do We Balance an Equation / Reaction?  Makes it look pretty  I do what I’m told  Has something to do with conservation  Of mass

29. Gas Has Mass  The balanced equation says I should  mix two hydrogens with one oxygen  And then the reaction should go well  Let’s See!

30. So if…….  If the space between molecules is so large  and volume of gas at equal pressure  and equal temperature have the same number of molecules How can we compare things that are not gases?

31. Mass  We would weigh it. Of course.  But reactions go by numbers of molecules and atoms  Not by mass, which is in grams  I wonder if there is a conversion factor?

32. How Can We Measure Atoms, the Darn Things are Too Small?  We can compare the mass of atoms  This was done back in the 1800’s  Hydrogen was found to be the smallest element  Everything was “relative” to Hydrogen  So we can compare elements by using hydrogen as a standard

33. Consider This About Atoms  Each atom of hydrogen has 1.0 amu and uranium has 238 amu.  Which has more atoms, one gram of H or one gram of U?  Calcium has 40. amu per atom and helium has 4 amu per atom.  Approximately how many times more atoms are in one gram of helium than in one gram of calcium?

34. Wow  So if we can compare masses of atoms to each other  Then we could weigh them to get the proper proportion of them in a reaction  Wouldn’t that be convenient?  If only there were a conversion factor!