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Week 13: Lectures 37 39 Lecture 37: W 11/16 Lecture 38: F 11/18 Lecture 39: M 11/28 Reading: BLB Ch 4.6; 10.5; 8.8; 5.3 5.7 Homework: BLB 10: 57; 5: 4, 17, 29, 37, 39, 53, 55, 83, 85; 8: 65a, 67ac, 72ab, 74, 92a; Supp Rxns: 12 18; 8: 12 14; 5: 1 7 Reminder: Angel Quiz 12 on Thur 11/17 ALEKS Objective 13 due on Tues 11/29 Jensen Office Hour: 501 Chemistry Building Tuesdays and Thursdays 10:30 11:30 am Thanksgiving break: Nov 21 25 Final Exam: Monday Dec 12 2:30 4:20 pm 40 questions. A 100% Score earns 58 points out of 200 course points (29% of final grade). Did you miss an exam? Download a Makeup Request Form from Angel > Exam Schedule Jensen Chem 110 Chap 5 Page: 1

Solution Reactions Basic skills: How to calculate formula weight (molar mass) How to do the following conversions: gram mole concentrations mole Titration: find concentration of unknown solution Can be used with: Acid-Base reactions (neutralization) Precipitation reactions Redox reactions Method: React solution of unknown concentration with solution of known concentration (standard solution) At the end point or equivalence point (using an indicator), the reaction is stoichiometrically complete At the end point, you know the exact moles of each reactant in solution from the balanced equation. Jensen Chem 110 Chap 5 Page: 2

Example: How many ml of a 0.827 M KOH solution is required for neutralization of a 35.00 ml sample of 0.737 M H 2 SO 4 solution? A. 35.0 ml B. 1.12 ml C. 25.8 ml D. 62.4 ml E. 39.3 ml Jensen Chem 110 Chap 5 Page: 3

Practice Example: One common component of antacids is Mg(OH) 2. If an upset stomach contains 125 ml of 0.115 M HCl, how many grams of Mg(OH) 2 is required to completely neutralize the acid? A. 0.144 g B. 0.286 g C. 0.419 g D. 0.525 g E. 0.884 g Jensen Chem 110 Chap 5 Page: 4

Gas Phase Reactions When chemical reactions involve gases, the balanced equation provides the number of moles of reactants and products. The ideal gas equation provides the link between number of moles and P, V, T of gases. Example: Air Bag How many liters of N 2 at 735 mm Hg and 26 C are produced from 126 g NaN 3 (sodium azide)? 2 NaN 3 (s) 2 Na(s) + 3 N 2 (g) Jensen Chem 110 Chap 5 Page: 5

Practice Example: Humans consume glucose to produce energy, and the products are CO 2 and H 2 O. C 6 H 12 O 6 (s) + O 2 (g) CO 2 (g) + H 2 O(l) [unbalanced] What volume of CO 2 is produced during the consumption of 4.65 g glucose at body temperature (37 ºC) and 1 atm? A. 3.94 L B. 1.97 L C. 0.657 L D. 3.47 L E. 7.88 L Jensen Chem 110 Chap 5 Page: 6

Thermochemistry KINETIC ENERGY POTENTIAL ENERGY Mechanical moving mass 1/2mv 2 mass in a place where force can act Electrical moving charge electrostatic:q 1 Q 2 /d Light Chemical Sound Nuclear Heat Gravitational photons molecules moving uniformly molecules moving randomly bonds binding energy mgh, where g=9.8 m/s 2 Chemical energy: Potential Energy associated with Bonding H H + H H + O=O O H H + H O H 3 bonds 4 bonds Jensen Chem 110 Chap 5 Page: 7

Energy can be converted between various forms, but total energy remains constant Law of Conservation of Energy First Law of Thermodynamics All energy lost by a system under observation must be gained by the surroundings (and vice versa) system: what you are interested in Surroundings: everything else During energy conversion, some heat is always produced Jensen Chem 110 Chap 5 Page: 8

Changes in Energy ΔE = E final E initial Both E and ΔE are state functions State functions: only depend on the current state (composition, T, P), does not depend on path or history ΔE for path 1 = ΔE for path 2 System energy Surroundings ΔE is Surroundings energy System ΔE is ΔE = q + w Jensen Chem 110 Chap 5 Page: 9

Energy & Enthalpy When changes occur at constant pressure: ΔE = q p + w expansion ~ negligible ΔH = q p H is enthalpy ΔH is the change in enthalpy Both H and ΔH are also state functions ΔH is the quantity of thermal energy (heat) absorbed or released by a system at constant pressure Examples of enthalpy: Energy transfers accompany physical changes; ΔH fusion (heat of fusion) ΔH vaporization (heat of vaporization) Energy transfers accompany chemical changes Enthalpy of reactions ΔH rxn Jensen Chem 110 Chap 5 Page: 10

Thermochemical Equation: A balanced chemical equation that also includes the enthalpy change ΔH = H products H reactants = ΔH rxn thermic reaction ΔH < 0 thermic reaction ΔH > 0 Characteristics of Enthalpy: 1) Enthalpy is an extensive property 2) ΔH rxn is equal in magnitude but opposite in sign for ΔH of reverse reaction 3) ΔH rxn depends on states of reactants & products (e.g., gas, liquid ) ΔH (delta H standard) standard P (1 atm) & T (usually 25 C) Note: NOT the same as STP for gases!!! Jensen Chem 110 Chap 5 Page: 11

Study the thermochemical equation of: The Hydrogen Balloon 2H 2 (g) + O 2 (g) 2H 2 O(g) ΔH = 483.6 kj 2 mol H 2 (g) reacts with 1 mol O 2 (g), produces 2 mol of H 2 O(g) and gives off 483.6 kj heat A. Is this reaction exothermic or endothermic? B. How much heat is given off per mole of O 2? C. How much heat is given off per mole of H 2? D. How much heat will be given off if 10.0 g of H 2 is consumed? E. What is ΔH for 2H 2 O(g) 2H 2 (g) + O 2 (g)? F. How much heat will be needed to convert 9.0 g of water into hydrogen and oxygen? Jensen Chem 110 Chap 5 Page: 12

Practice Example: How much heat is released when 25.0 g of sodium peroxide (Na 2 O 2 ) undergo this reaction? 2 Na 2 O 2(s) + 2 H 2 O (l) 4 NaOH (s) + O 2(g) H = 126 kj A. 20.2 kj B. 40.4 kj C. 67.5 kj D. 80.8 kj E. 126 kj Jensen Chem 110 Chap 5 Page: 13

Calorimetry: Experimental measure of heat flow (used to determine H rxn ) review molar heat capacity & specific heat Note: H 2 O is usually part of the surroundings q surr = C surr m ΔT Energy Conversion: q system = q surr Measure T for surroundings in a controlled environment (calorimeter) Jensen Chem 110 Chap 5 Page: 14

Coffee Cup Calorimeter (Constant Pressure calorimeter): Measure temperature change in solutions to calculate enthalpy change (heat lost or gained) q rxn for reactions under constant pressure q rxn = q soln = C soln m ΔT Bomb calorimeter (constant volume calorimetry): Heat evolved during combustion is absorbed by calorimeter contents causing a rise in water temperature q rxn = C cal ΔT C cal = heat capacity of bomb calorimeter NOTE: because bomb calorimetry is at constant volume (not constant pressure), heat transferred is ΔE, not ΔH Jensen Chem 110 Chap 5 Page: 15

Quantitative Calorimetry Example When a student mixes 50 ml of 1.0 M HCl and 50 ml of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0 C to 27.5 C. Calculate the enthalpy change for the reaction in kj per mol of HCl, assuming that the total volume of the solution is 100 ml, its density is 1.0 g/ml, and its specific heat is 4.18 J/(g K). A. 2.7 kj/mol B. 2.7 kj/mol C. 54.4 kj/mol D. 54.4. kj/mol E. 108 kj/mol Jensen Chem 110 Chap 5 Page: 16

Scratch paper: Jensen Chem 110 Chap 5 Page: 17

Hess s Law Enthalpy is an extensive property ΔH for a reaction is equal in magnitude and opposite in sign to ΔH for the reverse reaction ΔH for a reaction depends on the states of the reactants and products (gas, liquid, solid) Hessʼs Law: ΔH for an overall reaction is equal to the sum of the individual steps Consequence of ΔH being a state function C A B ΔH 1 ΔH rxn ΔH 2 B B C ΔH 2 ΔH 1 A A + B B + C A C ΔH rxn ΔH rxn = ΔH 1 + ΔH 2 Jensen Chem 110 Chap 5 Page: 18

Example: Given the following information A. H 2 (g) + F 2 (g) 2HF(g) ΔH A = 537kJ B. 2H 2 (g) + O 2 (g) 2 H 2 O(g) ΔH B = 572kJ Determine ΔH for the reaction: C. 2F 2 (g) + 2H 2 O(g) 4HF(g) + O 2 (g) ΔH C =? Idea: find combinations of reactions such that n A + m B = C then n ΔH A + m ΔH B = ΔH C Solve: x A: H 2 (g) + F 2 (g) HF(g) x B: H 2 O(g) H 2 (g) + O 2 (g) C : 2 F 2 (g) + 2 H 2 O(g) 4 HF(g) + O 2 (g) x A: ΔH 1 = x ΔH A = kj x B: ΔH 2 = x ΔH B = kj ΔH C = ΔH 1 + ΔH 2 = kj Jensen Chem 110 Chap 5 Page: 19

Practice Example: Given the following information: 2 SO 2 (g)+ O 2 (g) 2 SO 3 (g) ΔH A = 196kJ 2 S(s) + 3 O 2 (g) 2 SO 3 (g) ΔH B = 790kJ What is ΔH rxn for the following reaction? S(s) + O 2 (g) SO 2 (g) A. + 986 kj B. 986 kj C. 594 kj D. + 594 kj E. 297 kj Jensen Chem 110 Chap 5 Page: 20

Heat of formation ΔH f (enthalpy of formation): heat given off (or absorbed) when elements combine to form a compound combine Elements Compound ΔH f ΔH f (standard enthalpy of formation): heat given off (or absorbed) to form 1 mole compound when all elements are in their standard states. Definition of Standard State 1. P = 1 atm 2. T = 25 C (298K) 3. element is in its most stable state (gas/liquid/solid) For an element in its standard state: ΔH f = 0 NOTE: ΔH f values can be found in BLB Table 5.3 or Appendix C Two requirements for a reaction (at 25 C and 1 atm) to have ΔH rxn = ΔH f : 1. 2. Jensen Chem 110 Chap 5 Page: 21

Standard States of the elements (Phase under standard conditions of 298K and 1 atm) 1. Metals: all are solid at 298K and 1 atm except one (Hg) 2. Semi metals (metalloids): all are solids at 298K and 1 atm 3. Nonmetals at 298K and 1 atm A: Noble gases (Group 8): atomic gases B: Diatomics: H 2, N 2, O 2, Group 7 (F 2, Cl 2, Br 2, I 2 ) H 2, N 2, O 2, F 2, Cl 2, are gases Br 2 is a liquid I 2 is a solid C: all other non-metals are solids:c (graphite), S 8, P, Se Jensen Chem 110 Chap 5 Page: 22

Examples of ΔH f : 1. ΔH f for methanol, CH 3 OH (l) 1 atm _ C(graphite) + _ H 2 (g) + _ O 2 (g) 25 C _ CH 3 OH(l) Balance the equation for 1 mole of product Reactants all in standard state ΔH f = 238.6 kj/mol Note: reactants might not have whole numbers for coefficients & phase is important (l, s, g) 2. For which of the following reactions (at 25 C and 1 atm) is ΔH rxn = ΔH f? A. H 2 (g) + F 2 (g) 2HF (g) B. NO (g) + 1/2 O 2 (g) NO 2 (g) C. Pb (s) + Cl 2 (g) PbCl 2 (s) D. 2 Na (s) + N 2 (g) + 3 O 2 (g) 2 NaNO 3 (s) E. S (s) + O 3 (g) SO 3 (g) Jensen Chem 110 Chap 5 Page: 23

Standard Enthalpy of a Reaction: ΔH rxn Heat of reaction (ΔH rxn ) when all reactants and products are in the standard state. Obtain ΔH rxn from ΔH f (on data sheet): ΔH rxn = Σn ΔH f (products) Σm ΔH f (reactants) n, m are stoichiometric coefficients of each individual product and reactant, respectively This is an application of Hessʼs Law elements energy Σm ΔH f (reactants) reactants ΔH rxn Σn ΔH f (products) products ΔH rxn = Σn ΔH f (products) Σm ΔH f (reactants) Jensen Chem 110 Chap 5 Page: 24

Example: Sugar is broken down in the body to produce carbon dioxide and water. Using the information in the table, determine how much energy is produced by the controlled combustion of one mol of sugar in the body. C 12 H 22 O 11 (s) + 12 O 2 (g) 11 H 2 O(g) + 12 CO 2 (g) H 2 O(g) CO 2 (g) C 12 H 22 O 11 (s) H f (kj/mol) -241.8-393.5-2221.0 A. 1585.7 kj B. 2856.3 kj C. 5160.8 kj D. 5644.8 kj E. 9602.8 kj NOTE: ΔH f values can be found in BLB Table 5.3 or Appendix C Jensen Chem 110 Chap 5 Page: 25

Practice Example: The thermite reaction below is used for welding. What is the ΔH rxn for the reaction involving 1 mole of Al? 2 Al(s) + Fe 2 O 3 (s) Al 2 O 3 (s) + 2 Fe (s) ΔH f of Al 2 O 3 (s) = 1669.8 kj/mol ΔH f of Fe 2 O 3 (s) = 822.2 kj/mol A. + 847.6 kj B. 847.6 kj C. + 1895 kj D. 423.8 kj E. 2492 kj Jensen Chem 110 Chap 5 Page: 26

Where does the Energy Come From? The overall reaction has two steps: breaking the original bonds, and forming new ones. Bond breaking is always thermic Bond formation is always thermic Bond Enthalpies (bond energy) can be used to estimate ΔH rxn. Jensen Chem 110 Chap 5 Page: 27

Covalent Bond Length & Energy Bond length: distance between nuclei bond Bond energy (kj/mol) Bond length (pm) C C 348 154 C=C 614 134 C C 839 121 More electrons shared, shorter bond length Shorter bond length, stronger the bond Bond (dissociation) energy (D): enthalpy of bond breaking reaction in the gas phase. D > 0 (ΔH > 0) For diatomics, D is ΔH of one reaction: H H(g) 2H(g) D H-H = ΔH rxn = 436kJ/mol For polyatomics, D is an averaged quantity H O H(g) H O(g) + H(g) ΔH rxn = 494kJ/mol H O(g) H(g) + O(g) ΔH rxn = 424kJ/mol D H O = 463 kj/mol * (value obtained from averaging over many molecules) Exact values are known for some bonds, but the average values are not exact for any one case (unlike ΔH f ) Jensen Chem 110 Chap 5 Page: 28

Estimating ΔH rxn using Bond Energies ΔH rxn ΣnD broken ΣmD formed ; (n, m = # bonds) Energy is given off ( ) when bonds form. Example: Based on the data below, what is the estimated ΔH rxn for the combustion of 1 mo of propene (CH 2 =CH-CH 3 )? Bond D (kj/mol) Bond D (kj/mol) C-C 348 C=C 614 C=O 799 C H 413 O=O 495 O H 463 A. 3809 kj B. + 1905 kj C. 1905 kj D. + 3809 kj E. 2438 kj Jensen Chem 110 Chap 5 Page: 29

Scratch paper: Jensen Chem 110 Chap 5 Page: 30

Practice Example: The reaction below is used to produce methanol: CO(g) + 2H 2 (g) CH 3 OH H rxn = 128kJ Use the data below to calculate the approximate C O bond enthalpy. Bond D (kj/mol) Bond D (kj/mol) C-O 358 C=O 799 C H 413 O H 463 C O?? H H 436 A. 417 kj/mol B. 574 kj/mol C. 687 kj/mol D. 896 kj/mol E. 1060 kj/mol Jensen Chem 110 Chap 5 Page: 31

Scratch paper: Jensen Chem 110 Chap 5 Page: 32

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