Energy & Chemistry

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: 4 e- + O2 + 2 H2O ---> 4 OH- H2/O2 Fuel Cell Energy, page 288

Energy & Chemistry 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

Potential & Kinetic Energy Potential energy — energy a motionless body has by virtue of its position.

Potential Energyon 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 NaCl — composed of Na+ and Cl- ions.

Potential Energyon 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

Potential & Kinetic Energy Kinetic energy — energy of motion • Translation

Potential & Kinetic Energy Kinetic energy — energy of motion.

Internal Energy (E) • PE + KE = Internal energy (E or U) • Int. E of a chemical system depends on • number of particles • type of particles • temperature

Internal Energy (E) • PE + KE = Internal energy (E or U)

Internal Energy (E) • The higher the T the higher the internal energy • So, use changes in T (∆T) to monitor changes in E (∆E).

Thermodynamics • 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.

System and Surroundings • SYSTEM • The object under study • SURROUNDINGS • Everything outside the system

T(system) goes down T(surr) goes up Directionality of Heat Transfer • Heat always transfer from hotter object to cooler one. • EXOthermic: heat transfers from SYSTEM to SURROUNDINGS.

T(system) goes up T (surr) goes down Directionality of Heat Transfer • Heat always transfers from hotter object to cooler one. • ENDOthermic: heat transfers from SURROUNDINGSto theSYSTEM.

Energy & 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.

Energy Change in Chemical Processes PE of system dropped. KE increased. Therefore, you often feel a T increase.

James Joule 1818-1889 UNITS OF ENERGY 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

Which has the larger heat capacity? HEAT CAPACITY The heat required to raise an object’s T by 1 ˚C.

Specific Heat 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 Substance Spec. Heat (J/g•K) H2O 4.184 Ethylene glycol 2.39 Al 0.897 glass 0.84 Aluminum

Specific Heat 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?

heat gain/lose = q = (sp. ht.)(mass)(∆T) Specific Heat 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? where ∆T = Tfinal - Tinitial q = (0.897 J/g•K)(25.0 g)(37 - 310)K q = - 6120 J Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

Heat/Energy TransferNo Change in State q transferred = (sp. ht.)(mass)(∆T)

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?

Heat Transfer Because of conservation of energy, q(Fe) = –q(H2O) (heat out of Fe = heat into H2O) or q(Fe) + q(H2O) = 0 q(Fe) = (55.0 g)(Cp)(23.1 ˚C – 99.8 ˚C) q(Fe) = –4219 • Cp q(H2O) = (225 g)(4.184 J/K•g)(23.1 ˚C – 21.0 ˚C) q(H2O) = 1977 J q(Fe) + q(H2O) = –4219 Cp + 1977 = 0 Cp = 0.469 J/K•g

Heat Transfer with Change of State Changes of state involve energy (at constant T) Ice + 333 J/g (heat of fusion) -----> Liquid water q = (heat of fusion)(mass)

Heat Transfer and Changes of State Requires energy (heat). This is the reason a) you cool down after swimming b) you use water to put out a fire. Liquid ---> Vapor + energy

Heating/Cooling Curve for Water Note that T is constant as ice melts

+333 J/g +2260 J/g Heat & Changes of State What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100 oC? Heat of fusion of ice = 333 J/g Specific heat of water = 4.2 J/g•K Heat of vaporization = 2260 J/g

Heat & Changes of State What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100 oC? 1. To melt ice q = (500. g)(333 J/g) = 1.67 x 105 J 2. To raise water from 0 oC to 100 oC q = (500. g)(4.2 J/g•K)(100 - 0)K = 2.1 x 105 J 3. To evaporate water at 100 oC q = (500. g)(2260 J/g) = 1.13 x 106 J 4. Total heat energy = 1.51 x 106 J = 1510 kJ

Abba’s Refrigerator CCR, pages 232-233

Chemical Reactivity 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

Chemical Reactivity 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.

Heat Energy Transfer in a Physical Process CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Heat transfers from surroundings to system in endothermic process.

Heat Energy Transfer in a 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.

CO2 gas ∆E = E(final) - E(initial) = E(gas) - E(solid) CO2 solid Energy Level Diagram for Heat Energy Transfer

Heat Energy Transfer in Physical Change • Gas molecules have higher kinetic energy. • Also, WORKis done by the system in pushing aside the atmosphere. CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Two things have happened!

heat energy transferred work done by the system energy change FIRST LAW OF THERMODYNAMICS ∆E = q + w Energy is conserved!

heat transfer in (endothermic), +q heat transfer out (exothermic), -q w transfer in (+w) w transfer out (-w) SYSTEM ∆E = q + w

ENTHALPY Most chemical reactions occur at constant P, so Heat transferred at constant P = qp qp = ∆H where H = enthalpy and so ∆E = ∆H + w (and w is usually small) ∆H = heat transferred at constant P ≈ ∆E ∆H = change in heat content of the system ∆H = Hfinal - Hinitial

ENTHALPY ∆H = Hfinal - Hinitial If Hfinal > Hinitial then ∆H is positive Process is ENDOTHERMIC If Hfinal < Hinitial then ∆H is negative Process is EXOTHERMIC

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

USING ENTHALPY Making liquid H2O from H2 + O2 involves twoexothermic steps. H2 + O2 gas H2O vapor Liquid H2O

USING ENTHALPY 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.

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

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

∆H along one path = ∆H along another 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.

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

Enthalpy Values Depend on how the reaction is written and on phases of reactants and 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) ∆H˚ = -286 kJ

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

Energy & Chemistry