Honors Chemistry Unit 8 Thermochemistry Thermochemistry is the study of the changes in energy that accompany chemical reactions and physical changes. (Ch. 16) A. What is energy- energy is the capacity to do work or produce heat. Think of energy as something that an object possesses. Think of heat and work as ways that objects exchange energy. It exists in two basic forms. 1. Chemical Potential Energy- energy due to the composition of the substance and the type of bonding between the atoms. Stronger the bonds, than the more potential energy. 2. Kinetic Energy- energy of motion of the molecules. Equation: K.E. = ½ mv2. Greater the velocity, than the greater the K.E. Also, as you already know: πΎ. πΈ. ≈ π‘πππ. 3. All chemical and physical processes are governed by the Law of Conservation of Energy (1st Law of Thermodynamics). Which states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed. B. All chemical reactions involve changes in energy. 1. When we analyze energy changes, we need to focus on a well-defined part of the universe. a. System- the part that is singled out for study. In a reaction usually the reactants and products. b. Surroundings- everything with which the system can exchange energy. 2. Breaking of bonds absorbs energy; endothermic process a. Endothermic reactions absorb energy from their surroundings and transfer it to the system. b. The change (β) in energy is positive. 3. Formation of bonds releases energy; exothermic process a. Exothermic reactions release energy from the system to the surroundings. b. The change (β) in energy is negative. 4. Change in energy called “change in enthalpy”— βπ» a. when energy required to break bonds > energy released to form new bonds, +βπ» (endothermic) 1. products at a higher energy state than reactants (weaker bonds) 2. surroundings lose energy (cool down) b. when energy required to break bonds < energy released to form new bonds, –βπ» (exothermic) 1. products at a lower energy state than reactants (stronger bonds) 2. surroundings gain energy (heat up) 1 βH°rxn= Hproducts−Hreactants c. 5. C. Thermochemical equation a. Chemical equation with βπ» 1. listed to the right of equation 2. included as reactant (endothermic) or product (exothermic) b. βπ» can be used in dimensional analysis process. The magnitude is based upon the stoichiometric amounts of reactants and products as written. c. Units: SI unit of energy is the joule (J) 1. calorie: 1 cal = 4.184 J 2. Calorie, also called a nutritional calorie: 1 kilocal = 1 Cal 3. 1000 J = 1kJ Thermochemical changes of state- many processes besides chemical reactions absorb or release heat. 1. Phase changes for water that are endothermic. a. Energy absorbed from surroundings to break hydrogen bonds to go from a solid to a liquid is called heat of fusion, Hfusion a.k.a melting ο· H20(s) → H20(l) βHfusion= 6.01 kJ/mol b. 2. Energy absorbed from surroundings to break hydrogen bonds to go from a liquid to a gas is called heat of vaporization Hvaporization a.k.a evaporation. ο· H20(l) → H20(g) βHvaporization= 40.7 kJ/mol Phase changes that are exothermic. a. Energy released into surroundings to form hydrogen bonds when a gas changes into a liquid is called heat of condensation Hcondensation ο· H20(g) → H20(l) βHcondensation = −40.7 kJ/mol b. Energy released into the surroundings to form hydrogen bonds to go from a liquid to a solid is called heat of solidification Hsolidification a.k.a freezing ο· H20(l) → H20(s) βHsolidification= −6.01 kJ/mol D. Calorimetry- measurement of heat flow, the transfer of heat that occurs in chemical and physical processes can be measured using a calorimeter. Reactants are put in an insulated container filled with water, where heat is exchanged between reactants and water, but no heat is lost. 1. Specific Heat Capacity (Cs)- quantity of heat required to change the temperature of one gram of a substance by 1 K or (1°C). Units: π± πβπ² 2 π = πͺπ × π × βπ» ο· Cs= specifc heat capacity of substance; H2O = 4.18 J/g•K ο· q= heat flow ο· m= mass in grams ο· βπ = ππ − ππ 3. Types of calorimetry a. constant pressure- used for processes that take place at constant pressure: coffee-cup calorimeter. ο· assuming no heat escapes, heat released by the system is absorbed by the water. ππ π¦π π‘ππ = −ππ€ππ‘ππ 2. ο· E. π = πͺπ × π × βπ» Hess’s Law - it states that if you can add two or more thermochemical equations to produce a final equation, then the sum of the enthalpy changes of each individual reaction is the enthalpy change for the final reaction. - Rule #1 If the reaction is reversed, the sign of the βH changes too. - Rule #2 If the coefficients are multiplied by a factor, then the βH is multiplied by the same factor. - Rule #3 The state of matter (s, l, g, aq) of reactants and products, affects whether they cancel out or not. βπ»ππ£πππππ = βπ»1 + βπ»2 +. . . F. Standard Enthalpy of Formation (βH°f)- βH values for a reaction in which a compound is formed from its constituent elements in their standard states, 298K and 1atm. βH°rxn = ∑nβH°f (products) − ∑nβH°f (reactants) G. βH using bond energy (BE) data Bond Energies in (kJ/mol) Single H C N O S F Cl Br H 436 413 391 463 339 567 431 366 C 348 293 358 259 485 328 276 N 163 201 272 200 243 O 146 190 203 S 266 327 253 218 F 155 253 237 Cl 242 218 Br 193 I I 299 240 243 208 175 151 Multiple C=C 614 C=N 615 C=O 799 N=N 418 N=O 607 O=O 495 O=S 523 S=S 418 CοΊC 839 CοΊN 891 CοΊO 1072 NοΊN 941 3 1. energy needed to break a bond (i.e. C–H) in a diatomic, gaseous molecule, which contains the bond type ο· is approximately the same for any molecule ο· gaseous species ο· positive value (+ BE) for breaking bonds ο· forming bonds (– BE) βH = BEbreaking – BEforming threshold temperature (Tthreshold) when βG = 0 Tthreshold = βHo/βSo d. summary chart for determining when βG < 0 -βH +βH +βS βG < 0 for all T βG < 0 when T > βHo/βSo -βS βG < 0 when T < βHo/βSo βG > 0 for all T c. H. Second Law of Thermodynamics: predicting spontaneous change. 1. Spontaneous change is one that occurs without ongoing outside intervention. The sign of βH by itself does not predict spontaneity. a. Spontaneous processes with βH < 0 3 ο· 2Fe + O2βΆ Fe2O3(s) βπ»°ππ₯π = −826 ππ½ 2 b. 2. Spontaneity can be predicted by a function called Entropy (S). Entropy is a measure of the disorder or randomness of particles in a system also a measure of energy dispersal in a system. a. 2nd Law of Thermodynamics states that for any spontaneous process, the entropy of the universe increases. βππ’πππ£ > 0 b. Entropy determines the direction of chemical and physical change. A chemical system always proceeds in a direction that increases the entropy of the universe. 3. Entropy increases (βππ’πππ£ > 0) for the following: a. phase change: solid βΆ liquid βΆ gas b. diffusion, effusion c. dissolving of solids d. increase of temperature e. increase in number of moles of a gas during a rxn. f. atomic size within a group; increases entropy. g. molecular complexity within the same state; CH4(g), C2H6(g), C3H8(g) 4. I. ° ο· CH4 + 2O2 βΆ CO2 + 2H2O βπ»ππ₯π = −802 ππ½ ° ο· H2O(l) βΆ H2O(s) βπ»ππ₯π = −6.02 ππ½ Spontaneous processes with βH > 0 ° ο· H2O(l) βΆ H2O(g) βπ»ππ₯π = 40.7 ππ½ ° ο· H2O(s) βΆ H2O(l) βπ»ππ₯π = 6.02 ππ½ ° ο· NaCl(s) βΆ Na+(aq) + Cl-(aq) βπ»π πππ = 3.9 ππ½ βS°rxn = ∑nβS° (products) − ∑nβS°(reactants) Gibbs Free Energy (G): Overall Energy State (19.5 to 19.6) 1. spontaneous chemical and physical changes occur because the products are at a lower energy state and/or more disordered a. Gibbs free energy accounts for both βH and βS b. change in free energy: βG= βHo – TβSo 2. determine if thermodynamically favorable (βG < 0) a. lower potential energy (-βH)—chemical reactions b. greater disorder (+βS)—physical changes 4