1. 6.1 Nature of Energy

2. What is Energy? Energy is defined as the capacity to do work or to produce heat Types of energy Potential and Kinetic Energy

3. Law of conservation of Energy Energy must be conserved Reactions either absorb (endothermic) heat, or give off heat to the surroundings (exothermic) What is given of or absorbed by the reaction, must equal to what the surroundings absorbs or gives off.

4. What is Heat and Temperature Heat is a measure of energy that is transferred between two objects with two different temperatures. Temperature is a measure of the average kinetic energy of the particles of a substance.

5. Work, System and Surroundings Reactions (System) have the ability to do work on the surroundings and the surroundings can do work on the system We divide the universe into two parts, the system (Where the reaction is) and the surroundings (everything outside the reactions) Work (w) done on the surroundings is given a (-) value Work (w) on the system is given a (+) value Since work is kinetic energy, work is part of considering the next topic of internal energy E

6. Chemical Energy The system is what we are concerned about. The surroundings include everything else outside the system. Exothermic reactions are reactions where energy flows out of the system into the surroundings and therefore E is (-) Endothermic reactions are reactions where the system absorbs energy from the surroundings and E is (+) Any energy given off or absorb must be conserved

7. Internal Energy The first law of thermodynamics is written ∆E= q + w Where E is internal energy, q is heat and w is work. Since both work and heat are energy they must both be considered in the calculation of internal energy of a system. ∆E = Energy of System(final) – Energy of system(initial), change is energy of the system is the internal energy of a system. Remember when work is done on the system it is (+), when it is done on the surroundings it is (-), the same for heat. Examples from the book 6.1 and 6.2

9. Thermochemistry Work When a process occurs in an open container, commonly the only work done is a change in volume of a gas pushing on the surroundings (or being pushed on by the surroundings).

10. Thermochemistry Work We can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston. w = −P  V

11. Internal Energy The first law of thermodynamics is written ∆E= q + w Where E is internal energy, q is heat and w is work. Since both work and heat are energy they must both be considered in the calculation of internal energy of a system. ∆E = Energy of System(final) – Energy of system(initial), change is energy of the system is the internal energy of a system. Remember when work is done on the system it is (+), when it is done on the surroundings it is (-), the same for heat. Examples from the book 6.1 and 6.2

12. Thermochemistry Suggested Problems  Page 286 #’s 20-40 Odd Problems Due  20-40 Even

13. Thermochemistry Enthalpy If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system. Enthalpy is the internal energy plus the product of pressure and volume: H = E + PV

14. Thermochemistry Enthalpy When the system changes at constant pressure, the change in enthalpy,  H, is  H =  (E + PV) This can be written  H =  E + P  V

15. Thermochemistry Enthalpy Since  E = q + w and w = −P  V, we can substitute these into the enthalpy expression:  H =  E + P  V  H = (q+w) − w  H = q So, at constant pressure the change in enthalpy is the heat gained or lost.

16. Thermochemistry Endothermicity and Exothermicity A process is endothermic, then, when  H is positive.

17. Thermochemistry Endothermicity and Exothermicity A process is endothermic when  H is positive. A process is exothermic when  H is negative.

18. Thermochemistry Enthalpies of Reaction The change in enthalpy,  H, is the enthalpy of the products minus the enthalpy of the reactants:  H = H products − H reactants

19. Thermochemistry Enthalpies of Reaction This quantity,  H, is called the enthalpy of reaction, or the heat of reaction.

20. Thermochemistry The Truth about Enthalpy 1.Enthalpy is an extensive property. 2.  H for a reaction in the forward direction is equal in size, but opposite in sign, to  H for the reverse reaction. 3.  H for a reaction depends on the state of the products and the state of the reactants.

21. Thermochemistry Calorimetry Since we cannot know the exact enthalpy of the reactants and products, we measure  H through calorimetry, the measurement of heat flow.

22. Thermochemistry Heat Capacity and Specific Heat The amount of energy required to raise the temperature of a substance by 1 K (1  C) is its heat capacity. We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.

23. Thermochemistry Heat Capacity and Specific Heat Specific heat, then, is Specific heat = heat transferred mass  temperature change s = q m  Tm  T

24. Thermochemistry Constant Pressure Calorimetry By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.

25. Thermochemistry Constant Pressure Calorimetry Because the specific heat for water is well known (4.184 J/mol-K), we can measure  H for the reaction with this equation: q = m  s   T

26. Thermochemistry Bomb Calorimetry Reactions can be carried out in a sealed “bomb,” such as this one, and measure the heat absorbed by the water.

27. Thermochemistry Bomb Calorimetry Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy,  E, not  H. For most reactions, the difference is very small.

28. Thermochemistry Bomb Calorimetry

29. Thermochemistry Hess’s Law  H is well known for many reactions, and it is inconvenient to measure  H for every reaction in which we are interested. However, we can estimate  H using  H values that are published and the properties of enthalpy.

30. Thermochemistry Hess’s Law Hess’s law states that “If a reaction is carried out in a series of steps,  H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”

31. Thermochemistry Hess’s Law Because  H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products.

32. Thermochemistry Enthalpies of Formation An enthalpy of formation,  H f, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.

33. Thermochemistry Standard Enthalpies of Formation Standard enthalpies of formation,  H f, are measured under standard conditions (25°C and 1.00 atm pressure). 

34. Thermochemistry Calculation of  H Imagine this as occurring in 3 steps: C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l)

35. Thermochemistry Calculation of  H Imagine this as occurring in 3 steps: C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l)

36. Thermochemistry Calculation of  H Imagine this as occurring in 3 steps: C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l)

37. Thermochemistry C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l) C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) Calculation of  H The sum of these equations is:

38. Thermochemistry Calculation of  H We can use Hess’s law in this way:  H =  n  H f(products) -  m  H f(reactants) where n and m are the stoichiometric coefficients. 

39. Thermochemistry C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) Calculation of  H  H= [3(-393.5 kJ) + 4(-285.8 kJ)] - [1(-103.85 kJ) + 5(0 kJ)] = [(-1180.5 kJ) + (-1143.2 kJ)] - [(-103.85 kJ) + (0 kJ)] = (-2323.7 kJ) - (-103.85 kJ) = -2219.9 kJ

40. Thermochemistry Energy in Foods Most of the fuel in the food we eat comes from carbohydrates and fats.

41. Thermochemistry Fuels The vast majority of the energy consumed in this country comes from fossil fuels.