Atomic Mass

2. Atomic Mass • Atoms are so small, it is difficult to discuss how much they weigh in grams. • Use atomic mass units. • an atomic mass unit (amu) is one twelth the mass of a carbon-12 atom. • The decimal numbers on the table are atomic masses in amu.

4. They are not whole numbers • Because they are based on averages of atoms and of isotopes.

5. Total mass= 635482 u Average mass= total mass/10000=63.55 u

7. The Mole • The mole is a number. • A very large number, but still, just a number. • 6.022 x 1023 of anything is a mole • A large dozen. • The number of atoms in exactly 12 grams of carbon-12.

8. The Mole • 1 mol of atoms of any element weighs the number value of atomic weight in g. • 1 mol Cu weighs 63.55 g • 1 mol Ga weighs 69.723 g

9. Molar mass, M • Mass of 1 mole of a substance. • Often called molecular weight. • To determine the molar mass of an element, look on the table. • To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

10. Find the molar mass of • CH4 : CH4 → C + 4 H mCH4= mC+ mH MCH4= MC + 4×MH • Mg3P2 : M = 3×MMg + 2×MP • Ca(NO3)3 • Al2(Cr2O7)3 • CaSO4 · 2H2O

11. Number of moles, n Avogadro’s number Nav=6.02x1023 mol-1

12. Number of moles, number of atoms in 10 g Al? Number of moles, mass of 5.00x1020 atom Co?

13. Number of molecules of isopentyl acetate in 1 ug? • Number of carbon atoms?

14. Mass of CO32- in 4.86 mol CaCO3?

15. Percent Composition • Percent of each element a compound is composed of. • What is the percent composition of C2H5OH?

16. Suppose we have 1 mol C2H5OH • Sum of all percentages=100% • 100 g of C2H5OH contains 52.14 g C, 13.13 g H and 34.73 g O.

17. Find the percent compositions of the following compounds: • Al2(Cr2O7)3 • CaSO4 · 2H2O

18. From percent composition, you can determine the empirical formula. • Empirical Formula the lowest ratio of atoms in a molecule. A sample is 59.53% C, 5.38%H, 10.68%N, and 24.40%O, what is its empirical formula? E.F. C13H14N2O4

19. Empirical Formula and Molecular Formula H2O2 EF: HO Such compound doesn’t exist! MF = n × EF H2O2 = 2 (HO) MMF=n × MEF 34 = 2 × 17 2×MH + 2×MO = 2 (MH+MO) Alkenes: C2H4 , C3H8 , C4H8 , C5H10 … EF : CH2

20. E.F. CH2Cl M.F. C2H4Cl2

21. Elemental Analysis A substance is known to be composed of C, H and nitrogen. When 0.1156 g of this substance is reacted with oxygen, 0.1638 g CO2 and 0.1676 g H2O. Find the empirical formula.

22. Assumptions: • 1) All C in sample converted into CO2.

23. 2) All H in sample converted into H2O.

24. E.F. CH5N

25. Chemical Equations • Are sentences that describe what happens in a chemical reaction. • Reactants ® Products CH4 + O2® CO2 + H2O • Redistribution of bonds, atoms don’t change! • Equations should be balanced: Have the same number of each kind of atoms on both sides. • Law of conservation of mass. CH4 + 2O2® CO2 + 2 H2O Balancing by inspection.

26. Abbreviations • (s) ¯ • (g) ­ • (l) • (aq) • heat • D • catalyst

27. Practice • Ca(OH)2 + H3PO4® H2O + Ca3(PO4)2 • Cr + S8® Cr2S3 • KClO3(s) ® Cl2(g) + O2(g) • Solid iron(III) sulfide reacts with gaseous hydrogen chloride to form solid iron(III) chloride and hydrogen sulfide gas. • Fe2O3(s) + Al(s) ® Fe(s) + Al2O3(s)

28. Stoichiometric Calculations C3H8 + O2® CO2 + H2O • Balance the equation! C3H8 + 5 O2® 3 CO2 + 4 H2O • Language of moles, not grams! What mass of oxygen will react completely with 96.1 g of propane?

29. C3H8 + 5 O2® 3 CO2 + 4 H2O 5 2.18 ?

30. Grams of CO2 produced? C3H8 + 5 O2® 3 CO2 + 4 H2O 3 2.18 ?

31. Mass of CO2 absorbed by 1.00 kg LiOH(s)? 2 1 41.8 ?

32. Antiacids

33. 1 1 0.0119 ? 1 2 0.0172 ?

34. Examples • One way of producing O2(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 O2(g) are produced? • How many grams of potassium chloride? • How many grams of oxygen?

35. Examples • A piece of aluminum foil 5.11 in x 3.23 in x 0.0381 in is dissolved in excess HCl(aq). How many grams of H2(g) are produced? • How many grams of each reactant are needed to produce 15 grams of iron form the following reaction? Fe2O3(s) + Al(s) ® Fe(s) + Al2O3(s)

36. Examples • K2PtCl4(aq) + NH3(aq) ® Pt(NH3)2Cl2 (s)+ KCl(aq) • what mass of Pt(NH3)2Cl2 can be produced from 65 g of K2PtCl4 ? • How much KCl will be produced? • How much from 65 grams of NH3?

37. Limiting Reactant

38. 1 2 5 ?=10 mol NH3 15 (not available) 3 2 9 ?=6 mol NH3 H2 is LR.

39. 3 mol nitrogen 12 mol hydrogen 1 2 3 ?=6 mol NH3 3 2 12 ?=8 mol NH3 1 4 (not available) N2 is LR.

40. Limiting Reactant • Reactant that determines the amount of product formed. • The one you run out of first. • Makes the least product. • N2 + H2® NH3 • What mass of ammonia can be produced from a mixture of 100. g N2 and 500. g H2 ? • How much unreacted material remains?

41. 1 2 3.57 ?=7.14 mol NH3 3 2 248 ?=165.3 mol NH3

42. Reaction Yield • The theoretical yield is the amount you would make if everything went perfect. • The actual yield is what you make in the lab. Percent Yield

43. 68.5 kg CO is reacted with 8.60 kg H2. Calculate the percent yield if the actual yield was 3.57x104 g CH3OH.

44. 1 1 2445.6 ?=2445.6 mol CH3OH 2 1 4265.9 ?=2132.9 mol CH3OH

45. Examples • 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. 2 Al + 3 Br2®2 AlBr3

46. Examples • Years of experience have proven that the percent yield for the following reaction is 74.3% Hg + Br2® HgBr2 If 10.0 g of Hg and 9.00 g of Br2 are reacted, how much HgBr2 will be produced?

47. Examples • Commercial brass is an alloy of Cu and Zn. It reacts with HCl by the following reaction Zn(s) + 2HCl(aq) ® ZnCl2 (aq) + H2(g) Cu does not react. When 0.5065 g of brass is reacted with excess HCl, 0.0985 g of ZnCl2 are eventually isolated. What is the composition of the brass (i.e. Zn% and Cu%)?