3. Atomic Mass l Atoms are so small, it is difficult to discuss how much they weigh in grams l Use atomic mass units. l an atomic mass unit (amu) is one twelfth the mass of a carbon-12 atom l This gives us a basis for comparison l The decimal numbers on the table are atomic masses in amu

4. They are not whole numbers l Because they are based on averages of atoms and of isotopes. l can figure out the average atomic mass from the mass of the isotopes and their relative abundance. l add up the percent as decimals times the masses of the isotopes.

5. Examples l There are two isotopes of carbon 12 C with a mass of 12.00000 amu(98.892%), and 13 C with a mass of 13.00335 amu (1.108%)

6. The Mole l Makes the numbers on the table the mass of the average atom l Average atomic mass l Just atomic mass

7. The Mole l The mole is a number l a very large number, but still, just a number l 6.022 x 10 23 of anything is a mole l a large dozen l The number of atoms in exactly 12 grams of carbon-12

8. Molar mass l mass of 1 mole of a substance l often called molecular weight. l To determine the molar mass of an element, look on the table. l To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

9. Balancing equations CH 4 + O 2  CO 2 + H 2 O ReactantsProducts C11 O23 H42

10. Balancing equations CH 4 + O 2  CO 2 + 2 H 2 O ReactantsProducts C11 O23 H424

11. Balancing equations CH 4 + O 2  CO 2 + 2 H 2 O ReactantsProducts C11 O23 H424 4

12. Balancing equations CH 4 + 2O 2  CO 2 + 2 H 2 O ReactantsProducts C11 O23 H424 4 4

13. Abbreviations ( s ),  for product) ( g ),  for product) l ( aq ) l heat  l catalyst

14. Practice Ca(OH) 2 + H 3 PO 4  H 2 O + Ca 3 (PO 4 ) 2 KClO 3 ( s )  Cl( g ) + O 2 ( g ) l Solid iron(III) sulfide reacts with gaseous hydrogen chloride to form solid iron(III) chloride and dihydrogen monosulfide gas. Fe 2 O 3 ( s ) + Al( s )  Fe( s ) + Al 2 O 3 ( s )

15. Meaning l A balanced equation can be used to describe a reaction in molecules and atoms. l Not grams. l Chemical reactions happen molecules at a time l or dozens of molecules at a time l or moles of molecules.

16. Stoichiometry l Given an amount of either starting material or product, determining the other quantities. l use conversion factors from –molar mass (g - mole) –balanced equation (mole - mole) l keep track

17. Examples l One way of producing O 2 ( g ) involves the decomposition of potassium chlorate into potassium chloride and oxygen gas. A 25.5 g sample of Potassium chlorate is decomposed. How many moles of O 2 (g) are produced? l How many grams of potassium chloride? l How many grams of oxygen?

18. Example How many grams of each reactant are needed to produce 15 grams of iron from the following reaction? Fe 2 O 3 ( s ) + Al( s )  Fe( s ) + Al 2 O 3 ( s )

19. Examples K 2 PtCl 4 ( aq ) + NH 3 ( aq )  Pt(NH 3 ) 2 Cl 2 ( s )+ KCl( aq ) l what mass of Pt(NH 3 ) 2 Cl 2 can be produced from 65 g of K 2 PtCl 4 ? l How much KCl will be produced? l How much from 65 grams of NH 3 ?

20. Yield How much you get from an chemical reaction

21. Limiting Reagent l Reactant that determines the amount of product formed. l The one you run out of first. l Makes the least product. l Book shows you a ratio method. l It works. l So does mine

22. Limiting reagent l To determine the limiting reagent requires that you do two stoichiometry problems. l Figure out how much product each reactant makes. l The one that makes the least is the limiting reagent.

23. Example Ammonia is produced by the following reaction N 2 + H 2  NH 3 What mass of ammonia can be produced from a mixture of 500. g N 2 and 100. g H 2 ? l How much unreacted material remains?

24. Example l A 2.00 g sample of ammonia is mixed with 4.00 g of oxygen. Which is the limiting reactant and how much excess reactant remains after the reaction has stopped? l 4 NH 3 (g) + 5 O 2 (g)→4 NO(g) + 6 H 2 O(g)

25. Excess Reagent l The reactant you don’t run out of. l The amount of stuff you make is the yield. l The theoretical yield is the amount you would make if everything went perfect. l The actual yield is what you make in the lab.

26. Percent Yield l % yield = Actual x 100% Theoretical l % yield = what you got x 100% what you could have got

27. Examples l Aluminum burns in bromine producing aluminum bromide. In a laboratory 6.0 g of aluminum reacts with excess bromine. 50.3 g of aluminum bromide are produced. What are the three types of yield.

28. Examples Years of experience have proven that the percent yield for the following reaction is 74.3% Hg + Br 2  HgBr 2 If 10.0 g of Hg and 9.00 g of Br 2 are reacted, how much HgBr 2 will be produced? l If the reaction did go to completion, how much excess reagent would be left?

29. Examples Commercial brass is an alloy of Cu and Zn. It reacts with HCl by the following reaction Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g) Cu does not react. When 0.5065 g of brass is reacted with excess HCl, 0.0985 g of ZnCl 2 are eventually isolated. What is the composition of the brass?