1. General Chemistry M. R. Naimi-Jamal Faculty of Chemistry Iran University of Science & Technology

2. فصل پنجم: شیمی گرمایی

3. Contents 5-1Getting Started: Some Terminology 5-2Heat 5-3Heats of Reaction and Calorimetry 5-4Work 5-5The First Law of Thermodynamics 5-6Heats of Reaction:  U and  H 5-7The Indirect Determination of  H: Hess’s Law 5-8Standard Enthalpies of Formation 5-9Fuels as Sources of Energy

4. Thermochemistry is the study of energy changes that occur during chemical reactions. System: the part of the universe being studied. Surroundings: the rest of the universe. Thermochemistry: Basic Terms

6. Open: energy and matter can be exchanged with the surroundings. Closed: energy can be exchanged with the surroundings, matter cannot. Isolated: neither energy nor matter can be exchanged with the surroundings. Types of Systems A closed system; energy (not matter) can be exchanged.

7. Terminology Energy, U –The capacity to do work. Work –Force acting through a distance. Kinetic Energy –The energy of motion.

8. Energy Kinetic Energy ek =ek = 1 2 mv 2 [e k ] = kg m 2 s2s2 = J w = Fd [w ] = kg m s2s2 = J m Work

9. Energy Potential Energy –Energy due to condition, position, or composition. –Associated with forces of attraction or repulsion between objects. Energy can change from potential to kinetic.

10. Energy and Temperature Thermal Energy –Kinetic energy associated with random molecular motion. –In general proportional to temperature. –An intensive property. Heat and Work –q and w. –Energy changes.

11. Thermal Equilibrium Heat transfer ALWAYS occurs from a hotter object to a cooler object (directionality) Transfer of heat continues until both objects are at the same temperature (thermal equilibrium) The quantity of heat lost by a hotter object and the quantity of heat gained by a cooler object when they are in contact are numerically equal (required by the law of conservation of energy)

12. Units of Heat Calorie (cal) heat required to raise the temperature of 1.00 g of pure liquid water from 14.5 to 15.5 o C. Joule (J):SI unit for heat 1 cal = 4.184 J

13. Heat Capacity The quantity of heat required to change the temperature of a system by one degree. –Specific heat capacity, c. System is one gram of substance –Heat capacity, C. C = Mass x specific heat (mc). q = mcTmcT CTCT

14. Many metals have low specific heats. The specific heat of water is higher than that of almost any other substance.

15. Example: How much heat, in joules and in kilojoules, does it take to raise the temperature of 225 g of water from 25.0 to 100.0 ° C? Example: What will be the final temperature if a 5.00-g silver ring at 37.0 ° C gives off 25.0 J of heat to its surroundings? The specific heat of silver is 0.235 Jg -1 °C -1.

16. Example: An Estimation Example Without doing detailed calculations, determine which of the following is a likely approximate final temperature when 100 g of iron at 100 °C is added to 100 g of 20 °C water in a calorimeter, the specific heat of iron is 0.449 Jg -1 °C -1 : (a) 20 °C (b) 30 °C (c) 60 °C (d) 70 °C.

17. Conservation of Energy In interactions between a system and its surroundings the total energy remains constant. energy is neither created nor destroyed. q system + q surroundings = 0 q system = -q surroundings

18. Energy entering a system carries a positive sign: –heat absorbed by the system, or –work done on the system Energy leaving a system carries a negative sign –heat given off by the system –work done by the system First Law: Sign Convention

19. Heats of Reaction and Calorimetry Chemical energy. –Contributes to the internal energy of a system. Heat of reaction, q rxn. –The quantity of heat exchanged between a system and its surroundings when a chemical reaction occurs within the system, at constant temperature.

20. An exothermic reaction gives off heat –In an isolated system, the temperature increases. –The system goes from higher to lower energy; –q rxn is negative. An endothermic reaction absorbs heat –In an isolated system, the temperature decreases. –The system goes from lower to higher energy; –q rxn is positive. Heats of Reaction (q rxn )

23. Internal Energy Internal Energy, U. –Total energy (potential and kinetic) in a system. Translational kinetic energy. Molecular rotation. Bond vibration. Intermolecular attractions. Chemical bonds. Electrons.

24. First Law of Thermodynamics A system contains only internal energy. –A system does not contain heat or work. –These only occur during a change in the system. Law of Conservation of Energy –The energy of an isolated system is constant  U = q + w q = heat gained by system w = work gained (+) by or done on system (-)

25. Example: A gas does 135 J of work while expanding, and at the same time it absorbs 156 J of heat. What is the change in internal energy?  U = q + w = 156 – 135 = 21 J

26. The state of a system: its exact condition at a fixed instant. State is determined by the kinds and amounts of matter present, the structure of this matter at the molecular level, and the prevailing pressure and temperature. A state function is a property that has a unique value that depends only on the present state of a system, and does not depend on how the state was reached (does not depend on the history of the system). State Functions

27. Water at 293.15 K and 1.00 atm is in a specified state. d = 0.99820 g/mL This density is a unique function of the state. It does not matter how the state was established.

28. Functions of State U is a function of state. –Not easily measured.  U has a unique value between two states. –Is easily measured.

29. Path Dependent Functions For example: Fuel-Consume to go from one point to the another one is NOT a function of state, but a path dependent function

30. For a system where the reaction is carried out at constant volume,  V = 0 and  U = q V. Internal Energy Change at Constant Volume All the thermal energy produced by conversion from chemical energy is released as heat; no P-V work is done.

31. Work In addition to heat effects chemical reactions may also do work. Gas formed pushes against the atmosphere. Volume changes. Pressure-volume work.

32. Heats of Reaction:  U and  H Reactants → Products U i U f  U = U f - U i  U = q rxn + w In a system at constant volume: P  V=0  U = q rxn + 0 = q rxn = q v But we live in a constant pressure world! How does q p relate to q v ?

33. lnternal Energy Change at Constant Pressure For a system where the reaction is carried out at constant pressure,  U = q P – P  V or  U + P  V = q P Most of the thermal energy is released as heat. Some work is done to expand the system against the surroundings (push back the atmosphere).

34. Heats of Reaction

35.  U = q P + w We know that w = - P  V (minus because done by system) therefore:  U = q P - P  V q P =  U + P  V These are all state functions, so define a new function. Let H = U + PV Then  H = H f – H i =  U +  PV=  U + P  V + V  P If we work at constant pressure and temperature: V  P = 0  H =  U + P  V = q P

36. Standard States and Standard Enthalpy Changes Define a particular state as a standard state. Standard enthalpy of reaction,  H ° –The enthalpy change of a reaction in which all reactants and products are in their standard states. Standard State –The pure element or compound at a pressure of 1 atm and at the temperature of interest (usually 25 °C).

37. Enthalpy Diagrams

38. The enthalpy change that occurs in the formation of one mole of a substance in the standard state from the reference forms of the elements in their standard states. The standard enthalpy of formation of a pure element in its reference state is 0. Hf°Hf° Standard Enthalpies of Formation

39. When we say: “The standard enthalpy of formation of CH 3 OH(l) is 238.7 kJ”, we are saying that the reaction: C(graphite) + 2 H 2 (g) + ½ O 2 (g)  CH 3 OH(l) has a value of ΔH of –238.7 kJ. Standard Enthalpy of Formation We can treat ΔH f ° values as though they were absolute enthalpies, to determine enthalpy changes for reactions. Question: What is ΔHf° for an element in its standard state [such as O 2 (g)]? Hint: since the reactants are the same as the products …

40. Standard Enthalpies of Formation

41. Enthalpy of Reaction  H rxn = ∑  H f ° products - ∑  H f ° reactants

42. Table 7.3 Enthalpies of Formation of Ions in Aqueous Solutions

43. Example: Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of organic compounds. One reaction for producing synthesis gas is 3 CH 4 (g) + 2 H 2 O(l) + CO 2 (g)  4 CO(g) + 8 H 2 (g) ΔH° = ? Use standard enthalpies of formation from Table 6.2 to calculate the standard enthalpy change for this reaction. Example: The combustion of isopropyl alcohol, common rubbing alcohol, is represented by the equation 2 (CH 3 ) 2 CHOH(l) + 9 O 2 (g)  6 CO 2 (g) + 8 H 2 O(l) ΔH° = –4011 kJ Use this equation and data from Table 6.2 to establish the standard enthalpy of formation for isopropyl alcohol.

44. Fuels as Sources of Energy Fossil fuels. –Combustion is exothermic. –Non-renewable resource. –Environmental impact.