Energetics

1. Energetics

2. Thermochemistry and Thermodynamics Dr. Pál Bauer 2009

3. PartIThermochemistry Thermochemistry deals with the energy (heat) transfer during chemical reactions

4. Key Concepts of Lecture • Energy. • • Temperature. • • Zeroth law of thermodynamics. • • Internal energy. • • State functions. • • Transfer of energy in forms of heat and work. • • First law of thermodynamics. • • Enthalpy. • • Thermochemical equations. • .

5. • Calorimetry. • Hess’s law. • Standard enthalpy of formation. • Problems.

6. Energy • The ability to do work or transfer heat. • Work: Energy used to cause an object that has mass to move. • Heat: Energy used to cause the temperature of an object to rise. E = q + w

7. Energyis the ability to do work • Thermal energy is the energy associated with the random motion of atoms and molecules • Chemical energy is the energy stored within the bonds of chemical substances • Nuclear energy is the energy stored within the collection of neutrons and protons in the atom • Electrical energy is the energy associated with the flow of electrons • Potential energy is the energy available by virtue of an object’s position

8. Temperature = Thermal Energy 900C 400C Energy Changes in Chemical Reactions Heat is the transfer of thermal energy between two bodies that are at different temperatures. Temperature is a measure of the thermal energy. greater thermal energy 6.2

9. Temperature: Temperature is a measure of average atom/molecule speeds. For Example: At 25 oF, air molecules have an average speed ~1080 mi/hr At 100 oF, air molecules have an average speed ~1160 mi/hr

10. Temperature Temperature is a measure of average atom/molecule speeds. Temperature Measurement: Average molecule speeds are difficult to measure directly. Temperature is almost always measured indirectly (by one of it’s side effects). One such side effect is thermal expansion. Most materials expand when hotter and contract when cooler. Cool Iron bar

11. Temperature: Temperature is a measure of average atom/molecule speeds. Temperature Measurement: Average molecule speeds are difficult to measure directly. Temperature is almost always measured indirectly (by one of it’s side effects). One such side effect is thermal expansion. Most materials expand when hotter and contract when cooler. Hot Iron bar

12. Thermal expansion is a very small effect. For example: Steel expands 0.006% for every 10 oF From 20 oF to 100 oF: A 1.0 m steel bar expands 0.5 mm From 20 oF to 100 oF: A 500 ft steel bridge expands 3 in.

13. Thermal expansion is a very small effect. The effect is greatly exaggerated in a ‘liquid in glass’ thermometer.

14. Thermal expansion is a very small effect. Temperature Scales 212 oF 100 oC 32 oF 0 oC - Absolute temperature T(K) = tc + 273,15

15. Thermal expansion is a very small effect. Bi-metallic Strip or Spring Upon heating: metal 2 expands more than metal 1

16. The zeroth law of thermodynamics is a generalization about the thermal equilibrium among bodies, or thermodynamic systems, in contact. It results from the definition and properties of temperature. It can be stated as:"If A and C are each in thermal equilibrium with B, A is also in thermal equilibrium with C."

17. SURROUNDINGS SYSTEM Thermochemistry is the study of heat change in chemical reactions. The system is the specific part of the universe that is of interest in the study. closed isolated open energy nothing Exchange: mass & energy

18. Types of Energy • potential energy- (PE) due to position or composition • ex. attractive or repulsive forces • kinetic energy- (KE) due to motion of the object • KE = ½mv2 :depends on mass and volume

19. kg m2 1 J = 1  s2 Units of Energy • The SI unit of energy is the joule (J). • An older, non-SI unit is still in widespread use: The calorie (cal). 1 cal = 4.184 J

20. Transfer of Energy Etotal = Ek + Ep + Internal energy of system (U) • Two Ways to Transfer Energy: • Heat- (q) transfer of energy between two objects because of a temperature difference • Work- (w) force acting over a distance U = q + w ΔU = Δq + Δw

21. Internal Energy • (U) sum of potential and kinetic energy in system • can be changed by work, heat, or both • U = PE + KE • ∆U = q + w

22. Signs • signs are very important • signs will always reflect the system’s point of view unless otherwise stated

23. First Law of Thermodynamics • ΔU =Δq + Δw • Energy cannot be created nor destroyed. • Therefore, the total energy of the universe is a constant. • Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.

24. Pathway • the specific conditions of energy transfer • energy change is independent of pathway because it is a state function • work and heat depend on pathway so are not state functions • state function- depends only on current conditions, not past or future

25. Thermodynamics State functions are properties that are determined by the state of the system, regardless of how that condition was achieved. energy , pressure, volume, temperature Potential energy of hiker 1 and hiker 2 is the same even though they took different paths. 6.7

26. State Functions Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem.

27. State Functions • However, we do know that the internal energy of a system is independent of the path by which the system achieved that state. • In the system below, the water could have reached room temperature from either direction.

28. State Functions • Therefore, internal energy is a state function. • It depends only on the present state of the system, not on the path by which the system arrived at that state. • And so, E depends only on Einitial and Efinal.

29. State Functions • However, q and w are not state functions. • Whether the battery is shorted out or is discharged by running the fan, its E is the same. • But q and w are different in the two cases.

30. Transfer of Energy • exothermic- • energy is produced in reaction • flows out of system • container feels hot to the touch • endothermic- • energy is consumed by the reaction • flows into the system • container feels cold to the touch

31. Transfer of Energy Combustion of Methane Gas is exothermic

32. Transfer of Energy Reaction between nitrogen and oxygen is endothermic

33. Transfer of Energy • the energy comes from the potential energy difference between the reactants and products • energy produced (or absorbed) by reaction must equal the energy absorbed (or produced) by surroundings • usually the molecules with higher potential energy have weaker bonds than molecules with lower potential energy

34. 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).

35. 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.

36. Work • common types of work • expansion- work done by gas • compression- work done on a gas P is external pressure – not internal like we normally refer to

37. Example 1 • Find the ∆E for endothermic process where 15.6 kJ of heat flows in the system and 1.4 kJ of work is done on system • Since it is endothermic, q is + and w is +

38. Example 2 • Calculate the work of expansion of a gas from 46 L to 64 L at a constant pressure of 15 atm. • Since it is an expansion, ∆V is + and w is -

39. Example 3 • A balloon was inflated from 4.00 x 106 L to 4.50 x 106 L by the addition of 1.3 x 108 J of heat. Assuming the pressure is 1.0 atm, find the ∆E in Joules. (1 L∙atm=101.3 J) • Since it is an expansion, ∆V is + and w is -

40. 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 = U + PV

41. Enthalpy • When the system changes at constant pressure, the change in enthalpy, H, is H = (U + PV) • This can be written H = U + PV + VΔP

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

43. Endothermicity and Exothermicity • A process is endothermic, then, when H is positive.

44. Endothermicity and Exothermicity • A process is endothermic when H is positive. • A process is exothermic when H is negative.

45. Enthalpy of phase changesThe phase diagram of water

46. Phase transitions

47. Enthalpies of Reactions The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants: H = Hproducts−Hreactants

48. 2H2O (s) 2H2O (l) H2O (s) H2O (l) H2O (l) H2O (s) DH = -6.01 kJ DH = 6.01 kJ/mol ΔH = 6.01 kJ DH = 2 mol x 6.01 kJ/mol= 12.0 kJ Thermochemical Equations • The stoichiometric coefficients always refer to the number of moles of a substance • If you reverse a reaction, the sign of DH changes • If you multiply both sides of the equation by a factor n, then DH must change by the same factor n. 6.4

49. How much heat is evolved when 266 g of white phosphorus (P4) burn in air? x H2O (l) H2O (g) H2O (s) H2O (l) 3013 kJ 1 mol P4 x DH = 44.0 kJ DH = 6.01 kJ 1 mol P4 123.9 g P4 Thermochemical Equations • The physical states of all reactants and products must be specified in thermochemical equations. P4(s) + 5O2(g) P4O10(s)DHreaction = -3013 kJ = 6470 kJ 266 g P4 6.4

50. DH0 (O2) = 0 DH0 (O3) = 142 kJ/mol DH0 (C, graphite) = 0 DH0 (C, diamond) = 1.90 kJ/mol f f f f Standard enthalpy of formation (DH0) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm. f The standard enthalpy of formation of any element in its most stable form is zero. 6.6