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Chemical Reactions Week 12: Lectures 34 – 36 Lecture 34: W 11/9 Lecture 35: F 11/11 Lecture 36: M 11/14 Reading: BLB Ch 3.1 – 3.2; 3.6 – 3.7; 4.2 – 4.4 Homework: BLB 3: 1, 64; 4: 24, 39; Supp Rxns: 1 – 11 Reminder: No Angel Quiz on Thur 11/10 ALEKS Objective 12 due on Tues 11/15 Jensen Office Hour: 501 Chemistry Building Tuesdays and Thursdays 10:30 – 11:30 am Late drop deadline: Friday 11/11 @11:59 pm Final Exam: Monday Dec 12 2:30 – 4:20 pm !"#$%&'$("))$*+$,-*(.$!&/+0&*#$*$1*2,'3$ 4,5',)6$7&8($98&($:+;,0$<=$>-*($?@A,#'0,$$ Law of conservation of mass: total mass does not change during a chemical reaction • Mass of reactants MUST equal mass of products # of atoms of each element on reactant side = # of atoms of each element on product side Example: Complete combustion of pentane C5H12 + 8 O2 ! 5 CO2 + 6 H2O If mass of reactants is 100 g, then the mass of products must be ______ g. Reactants: ____C, ____H, ____O atoms Products: ____C, ____H, ____O atoms ___ mol C5H12 react completely with ____ mol O2, produce ____ mol CO2 and ____ mol H2O. Jensen Chem 110 Chap 3 & 4 Page: 2 Balancing Chemical Equations When reactants and products are both given, chemical reactions are balanced by changing the _____________. Four “Easy” Rules 1. Write the unbalanced molecular equation correctly (molecules involved with correct molecular formulas) C6H6 + O2 ! CO2 + H 2O Example: When the following equation C4H8O2(l) + O2(g) ! CO2(g) + H2O(g) is balanced with the smallest possible set of integer coefficients, the coefficient of O2 is A. B. C. D. E. 1 2 3 5 6 2. Balance the atoms of one element. C6H6 + O2 ! CO2 + H 2O 3. Balance atoms of remaining elements C6H6 + O2 ! CO2 + H 2O How many moles of water will be produced if 3 moles of C4H8O2 burn completely in air? 4. Check you work! Make sure that you use the smallest whole numbers. Jensen Chem 110 Chap 3 & 4 Page: 3 Jensen Chem 110 Chap 3 & 4 Page: 4 Patterns of Reactivity Patterns of Reactivity 1. Combination reactions (Chapter 3) 5. Exchange reactions (Chapter 4) (Double Displacement or Metathesis Rxn) Elements react to form compounds 2 Mg (s) + O2 (g) ! 2 MgO (s) Small compounds combine to form larger ones MgO (s) + CO2 (g) ! MgCO3 (s) Exchange reactions only occur if there is a driving force a. Precipitation Pb(NO3)2(aq) + 2 KI(aq) ! PbI2(s) " + 2 KNO3(aq) 2. Decomposition reactions (Chapter 3) 2 H2O (l) ! 2 H2 (g) + O2 (g) b.Neutralization (weak or non-electrolyte) CaCO3 (s) ! CaO (s) + CO2 (g) NaOH(aq) + HCl(aq) ! NaCl(aq) + H2O(!) 3. Complete Combustion reactions (Chapter 3) CH4 (g) + 2 O2 (g) ! CO2 (g) + 2 H2O(g) c. Gas formation 2 HCl(aq) + Na2S(aq) ! H2S(g) # + 2 NaCl(aq) All hydrocarbons will produce ___________ when How do you predict the phase of the products? they undergo complete combustion reactions. 4. Single displacement reactions (Chapter 4) Use Solubility Rules (Table 4.1) How do you know what is happening? Zn(s) + CuSO4(aq) ! ZnSO4(aq) + Cu(s) Jensen Chem 110 Chap 3 & 4 Page: 5 Use Net Ionic Equation Jensen Chem 110 Chap 3 & 4 Page: 6 Table 4.1: Aqueous Solubility of Ionic Compounds (Table provided on your data sheet) Soluble ionic compounds Important exceptions compounds containing NO3 – None C2H3O2 – None – Cl compounds of Ag+, Hg22+, and Pb2+ Br– compounds of Ag+, Hg22+, and Pb2+ – compounds of Ag+, Hg22+, and Pb2+ I SO42– compounds containing S Use solubility rules to predict reactions based on the reactants provided. ___________ drives the reaction. Mix MgCl2 and NaOH. What happens? compounds of Sr2+, Ba2+, Hg22+, and Pb2+ Insoluble ionic compounds Important exceptions 2– Demo: Precipitation reactions Mix KI and NaOH. What happens? compounds of the alkali metal cations and NH4+, Ca2+, Sr2+, and Ba2+ CO32– compounds of NH4+, the alkali metal cations PO4 3– compounds of NH4+, Mix AgNO3 and NaCl. What happens? the alkali metal cations OH – compounds of the alkali metal cations and NH4+, Ca2+, Sr2+, and Ba2+ Mix AgNO3 and KI. What happens? Simplified Rules (in WATER) 1. Almost all ammonium and alkali metal salts are soluble. 2. Most nitrates, acetates, chlorides, bromides, iodides, and sulfates are soluble. 3. Most sulfides, carbonates, phosphates, and hydroxides are insoluble. Jensen Chem 110 Chap 3 & 4 Page: 7 Jensen Chem 110 Chap 3 & 4 Page: 8 Net Ionic Equation: involves only the ions or molecules directly involved in the reaction 1st: Start with balanced molecular equation Pb(NO3)2(aq) + 2 KI(aq) ! PbI2(s) + 2 KNO3(aq) nd 2 : Dissociate all soluble strong electrolytes (strong acids, strong bases, soluble ionic salts) to get the Complete Ionic Equation. Subscripts for ions in the chemical formula (in the “Molecular” equation) become Coefficients for those ions in the Complete Ionic Equation. Practice Examples: 1. Which ions are spectator ions in the reaction represented by the following molecular equation? 2AgNO3(aq) + CaCl2(aq) !" 2AgCl(s) + Ca(NO3)2(aq) A. Ag+, Cl!, and Ca2+ B. Cl! and Ca2+ C. Ag+ and NO3! D. Ca2+ and NO3! E. Ca2+ 2. Mixing solutions of K2SO4(aq) and BaCl2(aq) produces an insoluble salt. What is the identity of the spectator ions? rd 3 : Identify spectator ions (ions that appear on both sides of equation) A. K+, SO42-, Ba2+, ClB. K+, SO42- 4th: Eliminate all spectator ions to get the Net Ionic Equation (ions and molecules that directly involved in the reaction) C. K+, Cl- Jensen Jensen Chem 110 Chap 3 & 4 Page: 9 D. Ba2+, ClE. Ba2+, SO42Chem 110 Chap 3 & 4 Page: 10 Examples of Acid–Base Neutralization Neutralization (Acid–Base) reactions acid + base ! salt + water • Acids: donate H+(aq) 1. Strong acid – strong base neutralization HCl (aq) + KOH (aq) ! KCl (aq) + H2O (l) # Complete ionic equation HCl (aq) ! • Bases: raises concentration of OH–(aq) ions # Spectator ions KOH (aq) ! • Salts: ionic compounds; replaces H+ of acid with positive ion e.g., HCl becomes KCl # Net ionic equation 2. Weak acid – strong base neutralization • Table 4.2—you"ve got this table memorized by now, right?? HF (aq) + KOH (aq) ! KF (aq) + H2O (l) # Complete ionic equation # Spectator ions # Net ionic equation Jensen Chem 110 Chap 3 & 4 Page: 11 Jensen Chem 110 Chap 3 & 4 Page: 12 Practice Example: Driving Force: Gas Formation What is the net ionic equation for the reaction between H3PO4 (aq) and KOH (aq)? • direct production of gas (e.g., H2, CO2, H2S) 2 HCl(aq) + Na2S(aq) ! H2S(g) # + 2 NaCl(aq) • production of weak acid which decomposes to a gas (e.g., H2CO3) A. H3PO4 (aq) + KOH (aq) ! K3PO4 (aq) + H2O (l) Example: Sodium Bicarbonate + HCl Molecular equation: B. H3PO4 (aq) + 3 KOH aq) ! K3PO4 (aq) + 3 H2O(l) C. H+ (aq) + OH— (aq) ! H2O (l) NaHCO3(aq)+HCl(aq) ! H2CO3(aq)+NaCl(aq) H2CO3(aq) ! CO2(g)+H2O(l) NaHCO3(aq)+HCl(aq)!CO2(g)+H2O(l)+NaCl(aq) D. H+(aq) + KOH (aq) ! K+ (aq) + H2O(l) Complete ionic equation Na+(aq) + HCO3$ (aq) + H+(aq) + Cl$ (aq) ! E. H3PO4 (aq) + 3 OH— (aq) Na+(aq) + Cl$ (aq) + CO2(g) + H2O(l) ! PO43— (aq) + 3 H2O(l) Net ionic equation: HCO3$ (aq) + H+(aq) ! CO2(g) + H2O(l) Spectator ions: Jensen Chem 110 Chap 3 & 4 Page: 13 Jensen Chem 110 Chap 3 & 4 Page: 14 Single Displacement Reactions (Oxidation-reduction; aka, redox reactions) Rules for determining Oxidation Numbers 1. The oxidation number of an atom of a pure Redox reactions: Reactions where electrons element is ______. are transferred from one reactant to another e.g. oxidation number for Cl2 is ___, for Fe is ___. 2. The oxidation number of a monatomic ion equals _____________. e.g. oxidation number for Cl– is ___, for Fe2+ is ___. 3. Some elements have the same oxidation number in almost all their compounds, and can be used as __________ to determine the oxidation numbers of • Reduction reaction: gaining e– X2 + 2 e – ! 2 X– e.g. Cl2 + 2 e– ! 2 Cl– other atoms in the compound. • Oxidation reaction: losing e– M ! Mn+ + n e– e.g. Fe ! Fe2+ + 2 e– 4. The sum of oxidation numbers in a neutral compound is ________; The sum of oxidation • Oxidation and reduction are always linked • Must be balanced: atoms/electrons/charge Jensen Chem 110 Chap 3 & 4 Page: 15 numbers in a polyatomic ion equals the __________ on the ion. Jensen Chem 110 Chap 3 & 4 Page: 16 Example: What is the oxidation number of Mn (Manganese) in MnO4–? Redox Reactions • Oxidizing reagents: Elements or compounds that oxidize the other reactant. e.g.: O2, halogens, H2O2, HNO3, Cr2O7–, MnO4– • Reducing agents: Elements or compounds that reduce the other reactant. e.g.: H2, C, metals Practice example: What is the oxidation state of S in H2SO4? A. B. C. D. E. +2 +4 +6 -2 -4 Jensen Oxidation numbers always change in redox reactions! Example: Balance the reaction between solid lead (II) oxide and ammonia gas to produce nitrogen gas, liquid water, and solid lead. Chem 110 Chap 3 & 4 Page: 17 Jensen Chem 110 Chap 3 & 4 Page: 18 Examples of Single Displacement Reactions Activity Series: predicts whether a certain metal will be oxidized by an acid or a salt 1. Metal + Salt Zn (s) + CuSO4 (aq) ! ZnSO4 (aq) + Cu (s) Ionic equation: Spectator Ions: Net ionic equation: What is the reducing agent (what is oxidized)? What is the oxidizing agent (what is reduced)? 2. Metal + Acid Zn (s) + 2 HCl (aq) ! ZnCl2 (aq) + H2 (g) What is the reducing agent (oxidized)? What is the oxidizing agent (reduced)? Jensen Chem 110 Chap 3 & 4 Page: 19 Jensen Chem 110 Chap 3 & 4 Page: 20 Using the Activity Series Problem Solving with Chemical Reactions Remember: An element that is __________ in the activity series will be oxidized by the ions of elements below it. Basic skills: • Avogadro"s number and the definition of mole • How to calculate formula weight (molar mass) • How to do the following conversions: gram % mole • Metals in the series will always be oxidized gram % molecules • What is meant by: by the ions of elements _______ empirical formula Zn(s) + AgNO3(aq) ! • Metals _______ H2 in series (e.g., Mg, Zn) will be oxidized by an acid (e.g., HCl) to form H2 Na(s) + H2O(l) ! • Metals toward bottom are unreactive (e.g., Ag, Pt, Au); that is, the elemental form is most stable Chem 110 Chap 3 & 4 We use these along with balanced chemical reactions to solve problems in chemistry & Use balanced chemical equation to connect: moles reactants % moles products & Use balanced chemical equation and conservation of mass to connect: grams reactants % grams products Example: C2H5OH + 3 O2 ! 2 CO2 + 3 H2O moles 1 mol 3 mol 2 mol 3 mol 1 mol of C2H5OH reacts with 3 mol of O2; produces 2 mol of CO2 and 3 mol of H2O Au(s) + H2O(g) ! Jensen molecular formula Page: 21 Jensen Chem 110 Chap 3 & 4 Page: 22 Steps to solve stoichiometry problems 1. Write the balanced chemical equation (or process) 2. Make a table & fill in given information Recognize: How much CO2 does your car produce? The combustion of octane (C8H18) in the presence of excess oxygen yields CO2 and H2O. If 2.6 kg of octane is consumed, how many kg of CO2 will it produce? • what you know already and what you are being asked for • what connections will take you from the knowns to unknowns 3. Make connections between measured properties and the balanced equation Mass molar mass moles Volume concentration moles (solutions) P, V, T (gases) Ideal Gas Law moles 4. Fill in table until you are able to solve the problem 5. Make sure your answer is REASONABLE Jensen Chem 110 Chap 3 & 4 Page: 23 Jensen Chem 110 Chap 3 & 4 Page: 24 Limiting Reagents & Reactant that is consumed completely; & Determines the final amount of product; & Must start with a balanced reaction. When reactants mixed in unbalanced proportions, some are left over (the ones in excess, unreacted) Be sure to test all reactants!!! • making a ham sandwich analogous to a chemical reaction Jensen Chem 110 Chap 3 & 4 Page: 25 Example: If 36.6 g of C2H5OH reacts with 63.8 g of O2 to form CO2 and H2O, how many grams of CO2 will be produced? A. B. C. D. E. 26.0g 43.2g 58.5 g 70.4 g 100.4g Jensen Chem 110 Chap 3 & 4 Page: 26 Practice Example: The combustion reaction Percent Yield Theoretical yield: the yield of product that results when the limiting reagent is completely consumed Actual yield: the yield you actually get in the real world Percent yield: ! actual yield $ % yield = # ' 100 " theoretical yield &% between 1.0 mole of C3H8 (g) and 1.0 mole of O2 (g) goes to completion: C3H8(g) + 5 O2(g) ! 3 CO2(g) + 4 H2O(g) Which of these statements are true? i. All of the C3H8 (g) is used up. ii. 3.0 moles of CO2 (g) is formed • Calculation is just one more step beyond a standard stoichiometry calculation • NOTE: if you get a % yield >100% something is wrong, you"ve just created matter!??!! Example continued: If you obtained 50.0 g of CO2 from your reaction, what is the percent yield? Jensen Chem 110 Chap 3 & 4 Page: 27 iii. 0.8 moles of H2O (g) is formed A. B. C. D. E. i only ii only i and ii only iii only i and iii only Jensen Chem 110 Chap 3 & 4 Page: 28 Practice Example: Lithium and nitrogen react to produce lithium nitride as follows: Scratch Paper: 6 Li (s) + N2 (g) ! 2 Li3N (s) If 5.00 g of each reactant undergo a reaction with a 80.5% yield, how many grams of Li3N are obtained from the reaction? A. B. C. D. E. 6.73 1.67 8.36 2.08 2.79 Jensen g g g g g Chem 110 Chap 3 & 4 Page: 29 Jensen Chem 110 Chap 3 & 4 Page: 30