Thermochemistry

1. Thermochemistry The study of energy and its transformations

2. Definitions When examining chemical systems or reactions, in order to keep track of energy changes, we consider the system and its surroundings. • The system is where we put our focus. Typically, it is the reactants and products. • The surroundings include everything else in the universe. Often, we just consider the immediate surroundings, such as the reaction vessel.

3. Energy Changes Chemical reactions and physical changes typically involve a transfer of energy. If a process such as melting ice requires energy, the reverse process of freezing water releases the same amount of energy.

4. Definitions If a reaction results in the evolution of heat, energy flows out of the system and into the surroundings. These reactions are exothermic. The energy lost by the system must be equal to the energy gained by the surroundings.

5. Definitions If a reaction requires heat, energy flows from the surroundings into the system. These reactions are endothermic. The energy gained by the system must be equal to the energy lost by the surroundings.

6. Physical Changes and Energy Changing the state of a substance involves an energy change. In melting or boiling a substance, the attractive forces between the atoms or molecules must be overcome, and heat is required. This process is endothermic. When a substance cools and condenses or freezes, heat is given off, and the process is exothermic.

7. Chemical Reactions and Energy In chemical reactions, the energy changes result from the breaking and the formation of chemical bonds. Bond breaking always requires energy. Bond making always releases energy.

8. Chemical Reactions and Energy Bond breaking always requires energy. Bond making always releases energy. In exothermic reactions, more energy is released in forming the products than is used in breaking apart the reactants.

9. Types of Systems Systems can be open (both energy and matter are exchanged, closed (only energy is exchanged) and isolated (neither energy nor matter is exchanged with the surroundings.)

10. Definitions - Energy • Chemical systems contain both kinetic energy and potential energy. Energy is the capacity to do work or to produce heat. An example of both is the combustion of gasoline. The gaseous products expand and do work (moving the pistons of an engine) and the reaction also produces heat.

11. Definitions - Energy Energy is the capacity to do work or to produce heat. Heat and work are ways that objects can exchange energy.

12. Definitions - Energy Kinetic energy is the energy of motion, and it depends upon the mass of the object and its velocity. Since molecules, especially those of gases, are in motion, they posess kinetic energy.

13. Definitions- Energy Potential energy is energy due to position or composition. Chemical energy is potential energy due to composition. For example, gasoline and oxygen have the potential to produce energy if they react.

14. Internal Energy The internal energy (E, or U) of a system is the sum of the kinetic and potential energies of all of the particles of the system. It is generally not possible to determine the internal energy of a system, but we can measure changes in internal energy. Internal energy is changed by the flow of work and/or heat.

15. Internal Energy Internal Energy is a state function. That is, it depends solely on the present state of the system, and not how it may have gotten to a particular state. A state function is independent of pathway.

16. Internal Energy Internal energy (E or U) is a state function, and depends only on the state of the system, and not how it got to that state.

17. Internal Energy & the 1st Law The First Law of Thermodynamics states that: Energy can be converted from one form to another, but cannot be created or destroyed. It is not possible to measure the total energy of a system, but it is possible to determine changes in energy.

18. Internal Energy Since energy may flow to or from the surroundings and the system, we are concerned with energy changes rather than the absolute value of the internal energy. ΔE = Efinal – Einitial (some texts use the symbol U for internal energy)

20. The 1st Law of Thermodynamics Since the energy change of the system is equal and opposite to the energy change of the surroundings, ΔEsystem= –ΔEsurroundings or ΔUsystem= –ΔUsurroundings

21. The 1st Law of Thermodynamics The implication of the first law of thermodynamics is that: The energy of the universe is constant.

22. Work and Heat Energy is the capacity to do work or transfer heat. The internal energy of a system changes when heat is exchanged, or when the system does work on its surroundings, or when the surroundings do work on the system. ΔE = q + w or ΔU = q + w

23. Energy, Work and Heat As the system loses energy to the surroundings, it can do so by losing heat (q), and/or doing work (w). As a result, ΔE = heat + work ΔE = q + w

24. Work Chemical reactions can be harnessed to do electrical work (as with batteries), or expansion work (such as expanding gases in an internal combustion engine). Our focus will be on expansion work.

25. Work Expansion work results when a reaction produces more gaseous products than reactants, and thus pushes back the atmosphere as the reaction proceeds. If the volume of the system contracts during reaction (more gaseous reactants than products), work is done by the surroundings on the system.

27. Expansion Work Usually we just consider the volumes of gases in a chemical reaction. 2 H2O(g)  2 H2(g) + O2(g) Since 2 moles of gaseous reactants produce 3 moles of gaseous products, the system expands, and does work in pushing back the atmosphere.

28. Expansion Work work = Force x Distance Pressure = Force/Area, or Force = Pressure(Area) work = Pressure(Area) x Distance work = Pressure(Area) x ∆h work = Pressure (length x width) x ∆h work = P ∆V

29. Expansion Work work = P ∆V If gases are produced by a reaction and the volume expands, the system is doing work on the surroundings. The sign, when considering the system, must be negative. So, work = - P ∆V

30. Energy, Work and Heat ΔE = q –PΔV Solving for q, the heat change, q = ΔE + PΔV The heat change, q, will vary with the reaction conditions. Reactions in an open vessel are performed at constant pressure, those in a sealed vessel are performed at constant volume.

31. Reactions at Constant Volume indicates constant volume q = ΔE + PΔV At constant volume, expansion work isn’t possible, and the heat change, qv, equals the change in internal energy. qv = ΔE or qv = ΔU

32. Energy, Work and Heat q = ΔE + PΔV At constant pressure, q becomes qp, and qp = ΔE + PΔV or qp = ΔU + PΔV denotes constant pressure

33. Energy, Work and Heat q = ΔE + PΔV At constant pressure, q becomes qp, and qp = ΔE + PΔV or qp = ΔU + PΔV denotes constant pressure

34. Energy and Enthalpy qp = ΔE + PΔV Since an open vessel is such a common apparatus, the heat transferred at constant pressure is given its own name, the enthalpy change, ΔH. qp = ΔE + PΔV = ΔH

35. Enthalpy Enthalpy is a state function, and is independent of reaction pathway. ΔH = Hfinal-Hinitial ΔH = Hproducts-Hreactants

36. Standard Conditions Many reactions are categorized by their standard enthalpy change, ΔHo. The degree sign indicates standard conditions. Standard conditions specify that the reactants and products are in the same molar amounts represented by the coefficients in the balanced chemical reaction.

37. Standard Conditions In a given experiment, the quantities of reactants and enthalpy change will vary, but the standard enthalpy change is reported based on molar quantities. The enthalpy of a reaction will also vary with the physical states of reactants or products as well as temperature and pressure.

38. Standard Conditions A thermodynamic standard state refers to a specific set of conditions. The standard is used so that values of enthalpy changes can be directly compared. The standard state is the most stable form of a substance at 1 atm pressure and 25oC.

39. Standard Conditions • Standard conditions are indicated using a degree symbol ( o ). Standard conditions for thermochemical data differ from the standard conditions used in the gas laws. 1. All gases have a pressure of exactly 1 atm. 2. Pure substances are in the form that they normally exist in at 25oC and 1 atm pressure. 3. All solutions have a concentration of exactly 1M.

40. Standard Conditions For example, since oxygen is a diatomic gas at 25oC, the standard state of oxygen is O2(g) at a pressure of 1 atm.

41. Thermochemical Equations Chemical reactions (or changes of state) may be written with their enthalpy of reaction. For example, CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l) ΔH=–890.4 kJ/mol The enthalpy change is for the reaction written, assuming molar quantities. 890.4 kJ of heat is produced when a mole of CH4(g) is burned.

42. Thermochemical Equations If the reaction is reversed, 890.4 kJ is required to produce each mole of CH4(g). CO2(g) + 2 H2O(l) CH4(g) + 2 O2(g) ΔH=+890.4 kJ/mol If the coefficients in a balanced reaction are multiplied by an integer, the value of ∆H is multiplied by the same integer.

43. Calorimetry Calorimetry is the science of measuring heat. It typically involves measuring temperature changes as a substance loses or gains heat. Since substances vary in how much their temperature changes as heat is lost or gained, it is important to know the heat capacity (C) of substances involved in the reaction.

44. Heat Capacity (C) The heat capacity of a substance is the amount of heat absorbed, usually in joules, per 1 degree (C or K) increase in temperature. The amount (mass) of the substance also determines the amount of heat lost or gained. C = heat absorbed = _q_ increase in temp. ΔT

45. Heat Capacity The specific heat capacity is for a gram of a substance. It has the units J/oC-g or J/K-g. The molar heat capacity is for a mole of a given substance. It has the units J/oC-mol or J/K-mol.

46. Calorimeter Constant In measuring heat changes during a reaction, any heat absorbed or lost by the calorimeter (the apparatus itself) must be considered. If this amount of heat is significant, the calorimeter constant may be provided or measured. This is the heat capacity of the specific apparatus used, and is expressed in J or kJ per degree change in temperature (K or oC).

47. Coffee Cup Calorimetry A simple device for determining heat changes of aqueous reactions at constant pressure is a coffee cup calorimeter. Since the contents are open to the atmosphere, the pressure, atmospheric pressure, remains constant during the reaction.

48. Coffee Cup Calorimetry The heat change for the reaction, qp, is equal to the enthalpy change for the reaction. If heat is given off, it goes towards warming up the contents of the calorimeter and toward warming up the calorimeter walls, thermometer, stirrer, etc.

49. Coffee Cup Calorimetry qreaction = qcontents + qcal qcontents = (mass of solution) (ΔTsoln)Csoln Csoln is the specific heat capacity of the reaction mixture. If solutions are aqueous and fairly dilute, the specific heat capacity of water, 4.18J/oC-g, may be used.

50. Coffee Cup Calorimetry qreaction = qcontents + qcal qcal = Ccal (ΔT) Ccal is the calorimeter heat capacity. It includes the heat needed to warm up the walls, thermometer and stirrer of the calorimeter, along with any heat loss due to leaks. In many simple calculations, Ccal is assumed to be negligible, and may be ignored.