Sheets Page 1 Lecture 35 Lecture 35: Reactions 2 Read: BLB 3.6–3.7 HW: BLB 3:57,64,73,79 Sup Rxns 1–11 Know: • oxidation/reduction reactions • stoichiometry calculations • limiting reagents • theoretical yield vs actual yield “NO$SCORE”$on$Exam$3??$talk$with$Mike$in$210$Whitmore$ Check out the grad-u-lator on the Chem110 website Review chemical nomenclature (e.g., BST #5, Lecture 8 …); & memorize those strong acids & bases (BLB Table 4.2) Bonus deadline for BST #10: Net ionic equations, Thurs, April 16; FINAL SKILL CHECK TEST DEADLINE: MONDAY, APRIL 27 Missed Exam 1, 2 or 3 due to illness? Make-up Exam: Monday, April 20 @ 6:30, 105 Wartik. Covers material from Exams 1–3, multiple choice. You must sign up by downloading & completing the request form & giving it to me: deadline to sign up is April 15. (info under “Exam Schedule” on Chem110 website) Late drop deadline: TODAY @ 11:59 pm via elion Need help?? Get help!! TAs in CRC (211 Whitmore) and SI—hours on Chem 110 website; my office hours (Mon 12:30-2 & Tues 10:30-12 in 324 Chem Bldg [or 326 Chem])Sheets Page 2 Lecture 35 Metal displacement reactions, or using the activity series of metals • activity series (BLB Table 4.5) shows relative ease of oxidation • metal on list can be oxidized by metal salts or acids below it Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s) • metals above H2 in series (e.g., Mg, Zn, Fe) will react with acid (e.g., HCl) to form H2 Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g) • metals toward bottom are unreactive (e.g., Ag, Pt, Au); that is, the elemental form is most stable—isnʼt it great that jewelry is made from these metals, rather than, say, K?Sheets Page 3 Lecture 35 Examples: Will the metal element displace the atom from its compound in a reaction? Na(s) + H2O(l) → Au(s) + H2O(g) → Zn(s) + AgNO3(aq) →Sheets Page 4 Lecture 35 Example: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) LEO goes GER What is oxidized? What is reduced? ionic equation: net ionic equation:Sheets Page 5 Lecture 35 Types of reactions (Table 3.1) • Combustion reactions organic molecule + O2 (g) → CO2 (g) + H2O (g) • Combination reactions A + B → C e.g., N2 (g) + 3 H2 (g) → 2 NH3 (g) • Decomposition reactions C → A + B e.g., C12H22O11 (s) → 12 C (s) + 11 H2O (g)Sheets Page 6 Lecture 35 To solve problems in chemistry… • write the balanced chemical reaction (or process) • make connections between experimentally measured properties and the balanced equation In other words… Given information such as mass, volume, pressure and temperature, how can one determine quantities of moles/molecules? Recognize (1) what you know already & what you are being asked for and (2) what connections will take you from knowns to unknownsSheets Page 7 Lecture 35 Using balanced chemical equations quantitiatively • use molar mass to convert: moles ⇔ grams • use balanced equation to convert: moles reactant ⇔ moles product • a balanced equation tells us the fewest number of molecules of each reactant it will take for the reaction to go to completion (i.e., NO reactants left), and the fewest number molecules of products that will then be formed 2 H2 + O2 → 2 H2O molecules: mol:Sheets Page 8 Lecture 35 Example: A How many g CO2 will be produced if we completely combust 12.00 g C6H12O6? B What is the mass of O2 consumed in this reaction? C How many g H2O are produced? C6H12O6+O2H2OCO2+MW180 g/mol 32 g/mol 44 g/mol 18 g/mol12.00 g ?????? ??? Does conservation of mass hold? That is, mass of reactants = mass of products??Sheets Page 9 Lecture 35 Example: 6.52 g of a compound containing C, H and O is completely combusted to yield 9.56 g CO2 and 3.91 g H2O. What is the empirical formula of the compound? CxHyOz+O2H2OCO2+MW32 g/mol 44 g/mol 18 g/mol6.52 g 9.56 g??? 3.91 g???emp formula?Sheets Page 10 Lecture 35 Before next class: Read: BLB 3.6–3.7 HW: BLB 3:57,64,73,79 Sup Rxns 1–11 Know: • stoichiometry calculations • limiting reagents • theoretical yield vs actual yield Answers: p. 8: (a) 17.6 g CO2; (b) 12.8 g O2; (c) 7.2 g H2O p. 9:
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