Energy & Chemistry

1. Energy & Chemistry 2H2(g) + O2(g) → 2H2O(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

2. 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 • 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.

3. Potential & Kinetic Energy Kinetic energy — energy of motion.

5. Internal Energy (E) • PE + KE = Internal energy (E or U) • Internal Energy of a chemical system depends on • number of particles • type of particles • temperature • The higher the T the higher the internal energy • So, use changes in T (∆T) to monitor changes in E (∆E).

6. Thermodynamics • Thermodynamics is the science of heat (energy) transfer. Heat transfers until thermal equilibrium is established. ∆T measures energy transferred. • SYSTEM • The object under study • SURROUNDINGS • Everything outside the system

7. 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.

8. 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.

9. 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 Potential Energy of system dropped. Kinetic energy increased. Therefore, you often feel a Temperature increase.

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

11. Specific heat capacity = heat lost or gained by substance (J) (mass, g) (T change, K) Specific Heat Capacity How much energy is transferred due to Temperature difference? The heat (q) “lost” or “gained” is related to a) sample mass b) change in T and c) specific heat capacity

12. Table of specific heat capacities Aluminum A Assuming an altitude of 194 meters above mean sea level (the world–wide median altitude of human habitation), an indoor temperature of 23 °C, a dewpoint of 9 °C (40.85% relative humidity), and 760 mm–Hg sea level–corrected barometric pressure (molar water vapor content = 1.16%).

13. 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?

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

15. 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?

16. Heating/Cooling Curve for Water Note that T is constant as ice melts or water boils

17. 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.

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

19. The First Law of Thermodynamics • Exothermic reactions generate specific amounts of heat. • This is because the potential energies of the products are lower than the potential energies of the reactants.

20. The First Law of Thermodynamics • There are two basic ideas of importance for thermodynamic systems. • Chemical systems tend toward a state of minimum potential energy. • Chemical systems tend toward a state of maximum disorder. • The first law is also known as the Law of Conservation of Energy. • Energy is neither created nor destroyed in chemical reactions and physical changes.

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

22. Some Thermodynamic Terms • Notice that the energy change in moving from the top to the bottom is independent of pathway but the work required may not be! • Some examples of state functions are: • T (temperature), P (pressure), V (volume), E (change in energy), H (change in enthalpy – the transfer of heat), and S (entropy) • Examples of non-state functions are: • n (moles), q (heat), w (work) ∆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.

23. Some Thermodynamic Terms • The properties of a system that depend only on the state of the system are called state functions. • State functions are always written using capital letters. • The value of a state function is independent of pathway. • An analog to a state function is the energy required to climb a mountain taking two different paths. • E1 = energy at the bottom of the mountain • E1 = mgh1 • E2 = energy at the top of the mountain • E2 = mgh2 • E = E2-E1 = mgh2 – mgh1 = mg(h)

24. Standard States and Standard Enthalpy Changes • Thermochemical standard state conditions • The thermochemical standard T = 298.15 K. • The thermochemical standard P = 1.0000 atm. • Be careful not to confuse these values with STP. • Thermochemical standard states of matter • For pure substances in their liquid or solid phase the standard state is the pure liquid or solid. • For gases the standard state is the gas at 1.00 atm of pressure. • For gaseous mixtures the partial pressure must be 1.00 atm. • For aqueous solutions the standard state is 1.00 M concentration. • ∆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

25. 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

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

27. 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

28. USING ENTHALPY Making liquid H2O from H2 + O2 involves twoexothermic steps. H2 + O2 gas H2O vapor Liquid H2O Making H2O from H2 involves two steps. H2(g) + 1/2 O2(g) → H2O(g) + 242 kJ H2O(g) → H2O(l) + 44 kJ H2(g) + 1/2 O2(g) → H2O(l) + 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.

29. 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 2H2(g) + O2(g) → 2H2O(g) ∆H˚ = -484 kJ H2O(g) → H2(g) + 1/2 O2(g) ∆H˚ = +242 kJ H2(g) + 1/2 O2(g) → H2O(l) ∆H˚ = -286 kJ

30. 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

31. Thermochemical equations are a balanced chemical reaction plus the H value for the reaction. For example, this is a thermochemical equation. The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles. 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6 mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions. Thermochemical Equations

32. Hess’s Law • Hess’s Law of Heat Summation, Hrxn = H1 +H2 +H3 + ..., states that the enthalpy change for a reaction is the same whether it occurs by one step or by any (hypothetical) series of steps. • Hess’s Law is true because H is a state function. • If we know the following Ho’s

33. Hess’s Law • For example, we can calculate the Ho for reaction [1] by properly adding (or subtracting) the Ho’s for reactions [2] and [3]. • Notice that reaction [1] has FeO and O2 as reactants and Fe2O3 as a product. • Arrange reactions [2] and [3] so that they also have FeO and O2 as reactants and Fe2O3 as a product. • Each reaction can be doubled, tripled, or multiplied by half, etc. • The Ho values are also doubled, tripled, etc. • If a reaction is reversed the sign of the Ho is changed.

34. Hess’s Law • Given the following equations and Hovalues calculate Ho for the reaction below.

35. Hess’s Law • Use a little algebra and Hess’s Law to get the appropriate Hovalues

36. This is an equivalent method of writing thermochemical equations. H < 0 designates an exothermic reaction. H > 0 designates an endothermic reaction Thermochemical Equations

37. Standard Molar Enthalpies of Formation, Hfo • The standard molar enthalpy of formation is defined as the enthalpy for the reaction in which one mole of a substance is formed from its constituent elements. • The symbol for standard molar enthalpy of formation is Hfo. • The standard molar enthalpy of formationfor MgCl2 is:

38. Standard Molar Enthalpies of Formation, Hfo • Standard molar enthalpies of formation have been determined for many substances and are tabulated in Table 15-1 and Appendix K in the text. • Standard molar enthalpies of elements in their most stable forms at 298.15 K and 1.000 atm are zero. • Example 15-4: The standard molar enthalpy of formation for phosphoric acid is -1281 kJ/mol. Write the equation for the reaction for whichHorxn = -1281 kJ. P in standard state is P4 Phosphoric acid in standard state is H3PO4(s)

39. Hess’s Law • Hess’s Law in a more useful form. • For any chemical reaction at standard conditions, the standard enthalpy change is the sum of the standard molar enthalpies of formation of the products (each multiplied by its coefficient in the balanced chemical equation) minus the corresponding sum for the reactants.

40. Hess’s Law

41. ∆Hfo, standard molar enthalpy of formation H2(g) + ½ O2(g) → H2O(g) ∆Hf˚ (H2O, g)= -241.8 kJ/mol C(s) + ½ O2(g) → CO(g)∆Hf˚ of CO = - 111 kJ/mol By definition, ∆Hfo= 0 for elements in their standard states. Use ∆H˚’s to calculate enthalpy change for H2O(g) + C(graphite) → H2(g) + CO(g)

42. Using Standard Enthalpy Values 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)

43. Standard Molar Enthalpies of Formation, Hfo • Calculate the enthalpy change for the reaction of one mole of H2(g) with one mole of F2(g) to form two moles of HF(g) at 25oC and one atmosphere.

44. Standard Molar Enthalpies of Formation, Hfo • Calculate the enthalpy change for the reaction in which 15.0 g of aluminum reacts with oxygen to form Al2O3 at 25oC and one atmosphere.

45. Hess’s Law • Calculate the H o298 forthe following reaction from data in Appendix K.

46. Hess’s Law • Application of Hess’s Law and more algebra allows us to calculate the Hfofor a substance participating in a reaction for which we know Hrxno , if we also know Hfofor all other substances in the reaction. • Given the following information, calculate Hfo for H2S(g).

47. Thermochemical Equations • Write the thermochemical equation for the reaction for CuSO4(aq) + 2NaOH(aq) Cu(OH)2(s) + Na2SO4(aq) 50.0mL of 0.400 M CuSO4 at 23.35 oC Tfinal 25.23oC 50.0mL of 0.600 M NaOH at 23.35 oC Density final solution = 1.02 g/mL CH2O = 4.184 J/goC