 Download Download Presentation 20 Sept. 2011

# 20 Sept. 2011

Download Presentation ## 20 Sept. 2011

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
##### Presentation Transcript

1. 20 Sept. 2011 • Objective: You will be able to: • Review making calculations and manipulating measurements • Determine the formula for copper (II) sulfate hydrate • Homework Quiz: Week of Sept. 19 • How will you know when to stop heating your copper (II) sulfate sample? Why?

2. Agenda • Turn in Problem Set 2 • Problem Set 1: Commonly missed questions • Finish lab procedure • Clean up Homework: Quiz on chapters 1-2: tomorrow Lab notebook due Thurs.

3. Clean Up • CuSO4 in trash • All equipment washed and hung up to dry • Ring stands, Bunsen burners and strikers, etc. back in cabinets • Wipe down counters • Start planning out your calculations

4. Homework • Quiz on ch. 1-2 tomorrow!

5. Safety • Wear goggles until all your equipment has been cleaned and returned. • A hot crucible looks just like a cold crucible! Always use crucible tongs. • Work efficiently but carefully.

6. Technical notes • Heat the crucible uncovered or with the cover tilted to allow water vapor to escape. • Cool the crucible with the cover on. • Cool the crucible in the desiccator for very best results. • Never mass a hot or warm crucible. • Oil from your fingers will stick to the crucible and effect your data.

7. Work Ethic • Work quickly. If you have “down time,” think: “What can I do now to save time later?” • Set up data tables and calculations while you wait.

8. This period • Carry out your procedure and collect data. • Begin calculations as soon as you can! • Percent of water in the hydrate by mass. • Mole ratio of anhydrous CuSO4 to H2O in your sample. • Work to show how you got your number of molecules of water of hydration.

9. Homework • Ch. 2 problem set: tomorrow • Quiz on chapters 1-2: tomorrow • Lab notebook due Thurs.

10. 21 September 2011 • Objective: You will be able to: • show what you know about chapters 1 and 2: calculations and compounds • Do now: Questions?

11. Quiz • Flip it over when you’re done. • Zero tolerance policy for disruptions. • When you finish: Work on lab calculations silently.

12. Homework • Check chapter 3 summer assignment answers online: Tomorrow • Lab notebook: calculations: tomorrow

13. 22 September 2011 • Take Out: Lab Notebook • Objective: You will be able to: • review Avogadro’s number, moles, molar mass, conversions, empirical and molecular formulas • Homework Quiz: • How many atoms of silver are equal to 3.50 moles of silver? • Calculate the molar mass of silver nitrate. (Nitrate is NO3-)

14. Agenda • Homework Quiz • Review of Avogadro’s number, moles, molar mass, conversions, empirical and molecular formulas • Challenge problem! Homework: p. 110 #18, 22, 26, 30, 44, 54: Mon.

15. The Mole • Atoms are so tiny; how do we quantify their masses?

16. The mole • Mole: The unit for amount of substance • =6.02x1023 atoms or molecules • Avogadro’s Number = 6.02x1023 • This is equal to the number of atoms of carbon in 12 grams of Carbon-12.

17. Example • How many moles of helium atoms are in 6.46 g of He? • How many atoms is this?

18. Problems • How many moles of magnesium are there in 87.3 g of Mg? • How many atoms is this? • Zinc is a silvery metal that is used in making brass (with copper). How many grams of Zn are in 0.356 mole of Zn?

19. Sulfur is a nonmetallic element present in coal. When coal is burned, sulfur is converted to sulfur dioxide and eventually to sulfuric acid that gives rise to the acid rain phenomenon. How many atoms are in 16.4 g of S? • Calculate the number of atoms in 0.551 g of potassium.

20. Molar Mass • The mass of a molecule. • Calculate the molar mass of: • sulfur dioxide (SO2) • caffeine (C8H10N4O2) • methanol (CH4O) SWBAT convert between grams, moles and number of particles of elements and compounds.

21. Using Molar Mass • Methane (CH4) is the principal component of natural gas. How many moles of methane are present in 6.07 g of CH4? • Calculate the number of moles of chloroform (CHCl3) in 198 g of chloroform.

22. Particles, Moles and Mass • Calculate the mass of 1.2x1024 molecules of carbon dioxide. • How many molecules of glucose are in 2.50x10-3 grams of glucose? • Calculate the mass of 1 molecule of water. • How many molecules of carbon dioxide are there in 0.0003 grams of carbon dioxide? SWBAT convert between grams, moles and number of particles of elements and compounds.

23. Percent Composition by Mass • the percent by mass of each element in a compound. • Can be used to determine the purity of a substance by comparing empirical data to known composition.

24. Practice Problems • Calculate the percent by mass of each element in hydrogen peroxide. • Phosphoric acid (H3PO4) is a colorless, syrupy liquid used in detergents, fertilizers, toothpastes and in carbonated beverages for a “tangy” flavor. Calculate the percent composition by mass of H, P and O in this compound.

25. Empirical Formulas • Given the percent composition, you can determine the empirical formula of a compound (reverse of calculating percent composition) • This data is often found by experiment.

26. Practice Problems • Ascorbic acid (vitamin C) cures scurvy. It is composed of 40.92% carbon, 4.58% hydrogen and 54.50% oxygen by mass. Determine its empirical formula. • A sample of a compound contains 1.52 grams of nitrogen and 3.47 grams of oxygen. The molar mass of this compound is between 90 and 95 g. Determine its molecular formula and accurate molar mass.

27. More practice problems • Determine the empirical formula of a compound having the following percent composition by mass: K: 24.75%, Mn: 34.77%, O: 40.51% • A sample of a compound containing boron and hydrogen contains 6.444g of boron and 1.803 g of hydrogen. The molar mass of the compound is about 30 g. What is its molecular formula and exact molar mass?

28. One more… • Chalcopyrite (CuFeS2) is a principal mineral of copper. Calculate the number of kilograms of copper in 3.71x103 kg of chalcopyrite.

29. Challenge! • A 0.1005 g sample of menthol (composed of C, H and O) is combusted, producing 0.2829 g of CO2 and 0.1159 g of H2O. What is the empirical formula? (Hint: Be aware that some oxygen in the product comes from the air, and some from the menthol!) • the compound has a molar mass of 156 g/mol what is the molecular formula?

30. Homework • p. 110 #18, 22, 26, 30, 44, 54: Mon. • Finish challenge problem

31. 26 September 2011 • Objective: You will be able to: • write and balance chemical equations and calculate mole to mole and mass to mass stoichiometry. • Homework Quiz • Calculate the mass of 1 atom of palladium. • Calculate the empirical formula of a compound composed of H, O and S, which has 2.1 percent H and 65.3 percent O.

32. Agenda • Homework Quiz • Go over homework • Review writing and balancing equations, mole and mass stoichiometry • Practice Problems Homework: 112 #60c, e, g, h, k and l, 64, 66, 71, 74: tomorrow Correct Quiz for a small quiz grade: Weds.

33. Decoding Chemical Equations • 2H2(g) + O2(g) → 2H2O(l)

34. Balancing Chemical Equations Ex 1. KClO3 KCl + O2 Ex 2. C2H6 + O2  CO2 + H2O

35. Write and balance: • Ammonia gas reacts with oxygen gas to produce nitrogen monoxide gas and liquid water. • Solid lithium reacts with nitrogen gas to produce solid lithium nitride. • Nitroglycerin (C3H5N3O9) decomposes explosively to produce nitrogen gas, carbon dioxide gas, water vapor and oxygen gas. • Carbon dioxide gas reacts with potassium hydroxide to produce potassium carbonate and water vapor.

36. Mole Ratios • Coefficients indicate the number of moles (or molecules) of each compound • Ex 1. N2(g) + 3H2(g)  2NH3(g) • How many moles of hydrogen gas react with 1 mole of nitrogen gas? • How many moles of ammonia gas are produced by the reaction of 3 moles of hydrogen gas in excess nitrogen gas?

37. N2(g) + 3H2(g)  2NH3(g) • How many moles of hydrogen gas are required to react with 15 moles of nitrogen gas? • How many moles of hydrogen gas are required to produce 25.0 moles of ammonia gas? • If 10 moles of nitrogen gas was reacted with 10 moles of hydrogen gas, which would be completely reacted? Which would be excess?

38. Mass to mass stoichiometric relationships: Ex 1. The food we eat is degraded in our bodies to provide energy for growth and function. A general equation for this very complex process is: C6H12O6 + 6O2 6CO2 + 6H2O If 856 g of C6H12O6 is consumed by a person over a certain period, what is the mass of CO2 produced?

39. Example 2 • All alkali metals react with water to produce hydrogen gas and the corresponding alkali metal hydroxide. A typical reaction is that between lithium and water: 2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g) How many grams of Li are needed to produce 9.89 grams of H2?

40. Example 3 • Methanol (CH3OH) burns in air according to the equation 2 CH3OH + 3O2 2CO2 + 4H2O If 209 g of methanol are used up in the combustion process, what is the mass of H2O produced?

41. Homework • 112 #60c, e, g, h, k and l, 64, 66, 71, 74

42. 27 September 2011 • Objective: You will be able to: • determine which reactant is the limiting reagent and calculate percent yield. • Homework Quiz • Balance: P4O10 + H2O → H3PO4 • If 10.0 grams of tetraphosphorusdecaoxide is reacted with excess water, calculate the number of grams of phosphoric acid produced. Bonus: What mass of water reacts with 10.0 grams of P4O10?

43. Agenda • Homework Quiz • Go over homework • Review limiting reagent and percent yield • Practice Problems Homework: p. 114 #84, 85, 89, 93, 107, 131 Correct Quiz for a small quiz grade: Weds. Ch. 1-3 test Thursday.

44. Example 4 • The reaction between nitric oxide and oxygen to form nitrogen dioxide is a key step in photochemical smog formation: 2NO(g) + O2(g)  2NO2(g) How many grams of O2 are needed to produce 2.21 g of NO2?

45. Limiting Reagents Ex 1. • Urea is prepared by reacting ammonia with carbon dioxide: 2NH3(g) + CO2(g) (NH2)2CO(aq) + H2O(l) In one process, 637.2 g of NH3 are treated with 1142 g of CO2. a) Which of the two reactants is the limiting reactant? b) Calculate the mass of (NH2)2CO formed. c) How much excess reagent (in grams) is left at the end of the reaction?

46. Example 2 • The reaction between aluminum and iron(III) oxide can generate temperatures approaching 3000oC and is used in welding metals: 2Al + Fe2O3 Al2O3 + 2Fe In one process, 124 g of Al are reacted with 601 g of Fe2O3. a) Calculate the mass (in grams) of Al2O3 formed. b) How much of the excess reagent is left at the end of the reaction?

47. 30 Sept. 2010 • Objective: SWBAT determine limiting reagent, and calculate percent yield. • Do now: TiCl4(g) + 2Mg(l)  Ti(s) + 2MgCl2(l) If 3.54x107 g of TiCl4 are reacted with 1.13x107 g of Mg, calculate the limiting reagent and the theoretical yield of Ti.