Chemistry and Chemical Reactivity 6th Edition PDF

Chemistry and Chemical Reactivity 1 6th Edition John C. Kotz Paul M. Treichel...

Chemistry and Chemical Reactivity 1 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 6 Principles of Reactivity: Energy and Chemical Reactions Lectures written by John Kotz © 2006 © 2006 Brooks/Cole Brooks/Cole Thomson - Thomson  2 © 2006 Brooks/Cole - Thomson  3 Geothermal power —Wairakei North Island, New Zealand © 2006 Brooks/Cole - Thomson  Energy & Chemistry 4 Burning peanuts Burning sugar supply sufficient (sugar reacts with energy to boil a cup KClO3, a strong of water. oxidizing agent) © 2006 Brooks/Cole - Thomson  Energy & Chemistry 5 These reactions are PRODUCT FAVORED They proceed almost completely from reactants to products, perhaps with some outside assistance. © 2006 Brooks/Cole - Thomson  6 Energy & Chemistry 2 H2(g) + O2(g) --> 2 H2O(g) + heat and light This can be set up to provide ELECTRIC ENERGY in a fuel cell. Oxidation: 2 H2 ---> 4 H+ + 4 e- Reduction: H2/O2 Fuel Cell 4 e- + O2 + 2 H2O ---> 4 OH- Energy, page 288 © 2006 Brooks/Cole - Thomson  Energy & Chemistry 7 ENERGY is the capacity to do work or transfer heat. HEAT is the form of energy that flows between 2 objects because of their difference in temperature. Other forms of energy — light electrical kinetic and potential © 2006 Brooks/Cole - Thomson  8 Potential & Kinetic Energy Potential energy — energy a motionless body has by virtue of its position. © 2006 Brooks/Cole - Thomson  Potential Energy 9 on the Atomic Scale Positive and negative particles (ions) attract one another. Two atoms can bond As the particles attract they have a lower potential NaCl — composed of energy Na+ and Cl- ions. © 2006 Brooks/Cole - Thomson  Potential Energy 10 on the Atomic Scale Positive and negative particles (ions) attract one another. Two atoms can bond As the particles attract they have a lower potential energy © 2006 Brooks/Cole - Thomson  11 Potential & Kinetic Energy Kinetic energy — energy of motion Translation © 2006 Brooks/Cole - Thomson  12 Potential & Kinetic Energy Kinetic energy — energy of motion. rotate vibrate translate © 2006 Brooks/Cole - Thomson  13 Internal Internal Energy Energy (E) (E) PE + KE = Internal energy (E or U) Int. E of a chemical system depends on number of particles type of particles temperature © 2006 Brooks/Cole - Thomson  14 Internal Internal Energy Energy (E) (E) PE + KE = Internal energy (E or U) © 2006 Brooks/Cole - Thomson  15 Internal Internal Energy Energy (E) (E) The higher the T the higher the internal energy So, use changes in T (∆T) to monitor changes in E (∆E). © 2006 Brooks/Cole - Thomson  Thermodynamics Thermodynamics 16 Thermodynamics is the science of heat (energy) transfer. Heat energy is associated with molecular motions. Heat transfers until thermal equilibrium is established. ∆T measures energy transferred. © 2006 Brooks/Cole - Thomson  System and Surroundings 17 SYSTEM – The object under study SURROUNDINGS – Everything outside the system © 2006 Brooks/Cole - Thomson  18 Directionality of Heat Transfer Heat always transfer from hotter object to cooler one. EXOthermic: heat transfers from SYSTEM to SURROUNDINGS. T(system) goes down T(surr) goes up © 2006 Brooks/Cole - Thomson  19 Directionality of Heat Transfer Heat always transfers from hotter object to cooler one. ENDOthermic: heat transfers from SURROUNDINGS to the SYSTEM. T(system) goes up T (surr) goes down © 2006 Brooks/Cole - Thomson  20 Energy Energy & & Chemistry Chemistry All of thermodynamics depends on the law of CONSERVATION OF ENERGY. The total energy is unchanged in a chemical reaction. If PE of products is less than reactants, the difference must be released as KE. © 2006 Brooks/Cole - Thomson  21 Energy Energy Change Change in in Chemical Chemical Processes Processes PE Reactants Kinetic Energy Products PE of system dropped. KE increased. Therefore, you often feel a T increase. © 2006 Brooks/Cole - Thomson  UNITS UNITS OF OF ENERGY ENERGY 22 1 calorie = heat required to raise temp. of 1.00 g of H2O by 1.0 oC. 1000 cal = 1 kilocalorie = 1 kcal 1 kcal = 1 Calorie (a food “calorie”) But we use the unit called the JOULE 1 cal = 4.184 joules James Joule 1818-1889 © 2006 Brooks/Cole - Thomson  23 HEAT CAPACITY The heat required to raise an object’s T by 1 ˚C. Which has the larger heat capacity? © 2006 Brooks/Cole - Thomson  Specific Specific Heat Heat Capacity 24 Capacity How much energy is transferred due to T difference? The heat (q) “lost” or “gained” is related to a) sample mass b) change in T and c) specific heat capacity Specific heat capacity = heat lost or gained by substance (J) (mass, g)(T change, K) © 2006 Brooks/Cole - Thomson  Specific Specific Heat Heat Capacity 25 Capacity Substance Spec. Heat (J/g K) H2O 4.184 Ethylene glycol 2.39 Al 0.897 glass 0.84 Aluminum © 2006 Brooks/Cole - Thomson  26 Specific Specific Heat Heat Capacity Capacity If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al? Specific heat capacity = heat lost or gained by substance (J) (mass, g)(T change, K) © 2006 Brooks/Cole - Thomson  27 Heat/Energy Transfer No Change in State q transferred = (sp. ht.)(mass)(∆T) © 2006 Brooks/Cole - Thomson  28 Heat Transfer Use heat transfer as a way to find specific heat capacity, Cp 55.0 g Fe at 99.8 ˚C Drop into 225 g water at 21.0 ˚C Water and metal come to 23.1 ˚C What is the specific heat capacity of the metal? © 2006 Brooks/Cole - Thomson  Heat Heat Transfer Transfer 29 with with Change Change of of State State Changes of state involve energy (at constant T) Ice + 333 J/g (heat of fusion) -----> Liquid water q = (heat of fusion)(mass) © 2006 Brooks/Cole - Thomson  30 Heat Heat Transfer Transfer and and Changes Changes of of State State Liquid ---> Vapor Requires energy (heat). This is the reason a) you cool down + energy after swimming b) you use water to put out a fire. © 2006 Brooks/Cole - Thomson  31 Heating/Cooling Heating/Cooling Curve Curve for for Water Water Note that T is constant as ice melts © 2006 Brooks/Cole - Thomson  Heat Heat & & Changes Changes of of State 32 State What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100 oC? Heat Heat of of fusion fusion of of ice ice == 333 333 J/g J/g Specific Specific heat heat of of water water == 4.2 4.2 J/g K J/g K Heat Heat of of vaporization vaporization == 2260 2260 J/g J/g +333 J/g +2260 J/g © 2006 Brooks/Cole - Thomson  33 Abba ’s Refrigerator Abba’s Refrigerator CCR, pages 232-233 © 2006 Brooks/Cole - Thomson  Chemical Reactivity 34 What drives chemical reactions? How do they occur? The first is answered by THERMODYNAMICS and the second by KINETICS. Have already seen a number of “driving forces” for reactions that are PRODUCT-FAVORED. formation of a precipitate gas formation H2O formation (acid-base reaction) electron transfer in a battery © 2006 Brooks/Cole - Thomson  Chemical Reactivity 35 But energy transfer also allows us to predict reactivity. In general, reactions that transfer energy to their surroundings are product- favored. So, let us consider heat transfer in chemical processes. © 2006 Brooks/Cole - Thomson  36 Heat Energy Transfer in a Physical Process CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Heat transfers from surroundings to system in endothermic process. © 2006 Brooks/Cole - Thomson  Heat Energy Transfer in a 37 Physical Process CO2 (s, -78 oC) ---> CO2 (g, -78 oC) A regular array of molecules in a solid -----> gas phase molecules. Gas molecules have higher kinetic energy. © 2006 Brooks/Cole - Thomson  38 Energy Level Diagram for Heat Energy Transfer CO2 gas ∆E = E(final) - E(initial) = E(gas) - E(solid) CO2 solid © 2006 Brooks/Cole - Thomson  39 Heat Energy Transfer in Physical Change CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Two things have happened! Gas molecules have higher kinetic energy. Also, WORK is done by the system in pushing aside the atmosphere. © 2006 Brooks/Cole - Thomson  40 FIRST LAW OF THERMODYNAMICS heat energy transferred ∆E = q + w work done by the energy system change Energy is conserved! © 2006 Brooks/Cole - Thomson  41 heat transfer in heat transfer out (endothermic), +q (exothermic), -q SYSTEM ∆E = q + w w transfer in w transfer out (+w) (-w) © 2006 Brooks/Cole - Thomson  42 ENTHALPY Most chemical reactions occur at constant P, so Heat transferred at constant P = qpp qpp = ∆H ∆H where H = enthalpy and so ∆ E=∆ ∆E H + w (and w is usually small) ∆H ∆ ∆HH = heat transferred at constant P ≈ ∆E ∆E ∆ ∆HH = change in heat content of the system ∆ ∆HH = Hfinal - Hinitial © 2006 Brooks/Cole - Thomson  43 ENTHALPY ∆H = Hfinal - Hinitial IfIfH Hfinal > > H H initial then then ∆H ∆H is is positive positive final initial Process is ENDOTHERMIC Process is ENDOTHERMIC IfIfH Hfinal < < H H initial then then ∆H ∆H is is negative negative final initial Process is EXOTHERMIC Process is EXOTHERMIC © 2006 Brooks/Cole - Thomson  44 USING ENTHALPY Consider the formation of water H2(g) + 1/2 O2(g) --> H2O(g) + 241.8 kJ Exothermic reaction — heat is a “product” and ∆H = – 241.8 kJ © 2006 Brooks/Cole - Thomson  USING ENTHALPY 45 Making liquid H2O from H2 + O2 involves two exothermic steps. H2 + O2 gas H2O vapor Liquid H2O © 2006 Brooks/Cole - Thomson  USING ENTHALPY 46 Making H2O from H2 involves two steps. H2(g) + 1/2 O2(g) ---> H2O(g) + 242 kJ H2O(g) ---> H2O(liq) + 44 kJ ----------------------------------------------------------------------- H2(g) + 1/2 O2(g) --> H2O(liq) + 286 kJ Example of HESS’S LAW— If a rxn. is the sum of 2 or more others, the net ∆H is the sum of the ∆H’s of the other rxns. © 2006 Brooks/Cole - Thomson  47 Hess’s Law & Energy Level Diagrams Forming H2O can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed. Active Figure 6.18 © 2006 Brooks/Cole - Thomson  48 Hess’s Law & Energy Level Diagrams Forming CO2 can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed. Active Figure 6.18 © 2006 Brooks/Cole - Thomson  49  ∆∆H H along along one one path path ==  ∆∆H H along along another another path path This equation is valid because ∆H is a STATE FUNCTION These depend only on the state of the system and not how it got there. V, T, P, energy — and your bank account! Unlike V, T, and P, one cannot measure absolute H. Can only measure ∆H. © 2006 Brooks/Cole - Thomson  50 Standard Enthalpy Values Most ∆H values are labeled ∆Ho Measured under standard conditions P = 1 bar Concentration = 1 mol/L T = usually 25 oC with all species in standard states e.g., C = graphite and O2 = gas © 2006 Brooks/Cole - Thomson  Enthalpy Values 51 Depend Depend onon how how the the reaction reaction is is written written and and on on phases phases of of reactants reactants and and products products H2(g) + 1/2 O2(g) --> H2O(g) ∆H˚ = -242 kJ 2 H2(g) + O2(g) --> 2 H2O(g) ∆H˚ = -484 kJ H2O(g) ---> H2(g) + 1/2 O2(g) ∆H˚ = +242 kJ H2(g) + 1/2 O2(g) --> H2O(liquid) © 2006 Brooks/Cole - Thomson ∆H˚ = -286 kJ  52 Standard Enthalpy Values NIST (Nat’l Institute for Standards and Technology) gives values of ∆Hfo = standard molar enthalpy of formation — the enthalpy change when 1 mol of compound is formed from elements under standard conditions. See Table 6.2 and Appendix L © 2006 Brooks/Cole - Thomson  53 ∆Hfo, standard molar enthalpy of formation H2(g) + 1/2 O2(g) --> H2O(g) ∆Hfo (H2O, g)= -241.8 kJ/mol By definition, ∆Hfo = 0 for elements in their standard states. © 2006 Brooks/Cole - Thomson  Using Standard Enthalpy Values 54 Use ∆H˚’s to calculate enthalpy change for H2O(g) + C(graphite) --> H2(g) + CO(g) (product is called “water gas”) © 2006 Brooks/Cole - Thomson  Using Standard Enthalpy Values 55 H2O(g) + C(graphite) --> H2(g) + CO(g) From reference books we find H2(g) + 1/2 O2(g) --> H2O(g) ∆Hf˚ of H2O vapor = - 242 kJ/mol C(s) + 1/2 O2(g) --> CO(g) ∆Hf˚ of CO = - 111 kJ/mol © 2006 Brooks/Cole - Thomson  Using Standard Enthalpy Values 56 H2O(g) --> H2(g) + 1/2 O2(g) ∆Ho = +242 kJ C(s) + 1/2 O2(g) --> CO(g) ∆Ho = -111 kJ -------------------------------------------------------------------------------- H22O(g) + C(graphite) -- > H22(g) + CO(g) --> ∆ ∆HHoonet net = +131 kJ To convert 1 mol of water to 1 mol each of H2 and CO requires 131 kJ of energy. The “water gas” reaction is ENDOthermic. © 2006 Brooks/Cole - Thomson  Using Standard Enthalpy Values 57 In general, when ALL Calculate ∆H of enthalpies of formation are reaction? known, ∆H oo = ∆H rxnrxn ∆∆HHffoo (products) - ∆ H ∆Hffoo (reactants) Remember that ∆ always = final – initial © 2006 Brooks/Cole - Thomson  Using Standard Enthalpy Values 58 Calculate the heat of combustion of methanol, i.e., ∆Horxn for CH3OH(g) + 3/2 O2(g) --> CO2(g) + 2 H2O(g) ∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react) © 2006 Brooks/Cole - Thomson  Using Standard Enthalpy Values 59 CH3OH(g) + 3/2 O2(g) --> CO2(g) + 2 H2O(g) ∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react) ∆Horxn = ∆Hfo (CO2) + 2 ∆Hfo (H2O) - {3/2 ∆Hfo (O2) + ∆Hfo (CH3OH)} = (-393.5 kJ) + 2 (-241.8 kJ) - {0 + (-201.5 kJ)} ∆Horxn = -675.6 kJ per mol of methanol © 2006 Brooks/Cole - Thomson  CALORIMETRY 60 Measuring Heats of Reaction © 2006 Brooks/Cole - Thomson  CALORIMETRY 61 Measuring Heats of Reaction Constant Volume “Bomb” Calorimeter Burn combustible sample. Measure heat evolved in a reaction. Derive ∆E for reaction. © 2006 Brooks/Cole - Thomson  62 Calorimetry Some heat from reaction warms water qwater = (sp. ht.)(water mass)(∆T) Some heat from reaction warms “bomb” qbomb = (heat capacity, J/K)(∆T) Total heat evolved = qtotal = qwater + qbomb © 2006 Brooks/Cole - Thomson  63 Measuring Measuring Heats Heats of of Reaction Reaction CALORIMETRY CALORIMETRY Calculate heat of combustion of octane. C8H18 + 25/2 O2 --> 8 CO2 + 9 H2O Burn 1.00 g of octane Temp rises from 25.00 to 33.20 oC Calorimeter contains 1200 g water Heat capacity of bomb = 837 J/K © 2006 Brooks/Cole - Thomson