THERMOCHEMISTRY- A

1. THERMOCHEMISTRY- A By Dr. Hisham E Abdellatef http://www.staff.zu.edu.eg/ezzat_hisham/browseMyFiles.asp?path=./userdownloads/physical%20chemistry%20for%20clinical%20pharmacy/

2. THERMOCHEMISTRY The study of heat released or required by chemical reactions or heat changes caused by chemical reactions

3. Energy Kinetic energy (EK) Potential energy (EP) Energy due to motion Energy due to position (stored energy) What is Energy?

4. Internal Energy (E) • Internal Energy, E, The total energy of a system • E(system) = EK (system) + EP (system)

5. Potential & Kinetic Energy Potential energy — energy a motionless body has by virtue of its position.

6. Potential & Kinetic Energy Kinetic energy Energy of Motion EK = ½ mv2 • i.e. the kinetic energy of an object depends on • both its mass and its speed.

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

8. Problem • What is the kinetic energy of a person whose mass is 130 lb (59.0 kg) traveling in a car at 60 mph (26.8 m/s). • The SI unit of energy, kg.m2/s2, is the Joule.

9. James Joule 1818-1889 James Joule 1818-1889 UNITS OF ENERGY 1 calorie = heat required to raise temp. of 1.00 g of H2O by 1.0 oC. 1000 cal = 1 kilocalorie = 1 kcal But we use the unit called the JOULE 1 cal = 4.184 joules

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

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

12. 1st Law of Thermodynamics: • Energy can neither created not destroyed, only transformed from one form to another

13. System and Surroundings In chemical reactions, heat is often transferred from the “system” to its “surroundings,” or vice versa. • The substance or mixture of substances under study in which a change occurs is called thethermodynamic system (or simply thesystem.) • Thesurroundingsare everything outside of the thermodynamic system.

14. SURROUNDINGS HEAT HEAT HEAT HEAT SYSTEM SYSTEM EXOTHERMIC ENDOTHERMIC

15. Burning fossil fuels is an exothermic reaction Exothermic If E < 0, Efinal < Einitial cellular respiration of glucose

16. Endothermic If E > 0, Efinal > Einitial Photosynthesis is an endothermic reaction (requires energy input from sun)

17. Heat of Reaction • An exothermic process is a chemical reaction or physical change in which heat is evolved (q is negative). • An endothermic process is a chemical reaction or physical change in which heat is absorbed (q is positive).

18. Exothermic Endothermic 6.2

19. Enthalpy Diagrams • Values of DH are measured experimentally. • Negative values indicate exothermic reactions. • Positive values indicate endothermic reactions. An increase in enthalpy during the reaction; DH is positive. A decrease in enthalpy during the reaction; DH is negative.

20. ExothermicExamples • Oxidation – wooden splint burning (giving off light, heat, CO2, H2O • Burning H2 in air, • body reactions, • dissolving metals in acid, • mixing acid and water, • sugar dehydration

21. Endothermic Examples • Electrolysis (breaking water down into H2 and O2 by running electricity in it) • Photosynthesis, pasteurization, canning vegetables • 2 H2 + O2 2H2O + energy • 4 g + 32 g  36 g 136 600 cal • 2H2O + energy  2H2 + O2 • 36 g 136 600 cal  4g + 32 g

22. Changes in Internal Energy • If E > 0, Efinal > Einitial • Therefore, the system absorbed energy from the surroundings. • This energy change is called endergonic.

23. Changes in Internal Energy • If E < 0, Efinal < Einitial • Therefore, the system released energy to the surroundings. • This energy change is called exergonic.

24. Enthalpy (H) of the reaction(Comes from Greek for “heat inside”) • the sum of internal energy and the product of this pressure and volume. H = E + PV • E is the internal energy, • P is the pressure and • V is the volume of the system. • It is also called heat content. • ΔH = H product – H reactants = Hp – Hr

25. Calculation of ΔH from ΔE • When the system changes at constant pressure, the change in enthalpy, H, is H = (E + PV) H = E + PV At constant pressure and temperature The enthalpy of a chemical is measured in kilojoules per mole (kJmol-1).

26. At constant volume H = E + PV For solid and liquid H = E

27. In case of gases H = E + PV (I) ΔV =Δn x V • Δn = no of moles of products - no of moles of reactants PxΔV = PVxΔn (II) • But PV = RT (for one mole of gas) • Putting RT in place of PV in equation (II) we get PΔV = RTΔn • Substituting the value of P AV in equation (I) we get ΔH = ΔE + Δn RT

28. R = 1.987 cal. (=2 cal.) or = 8.314 joules

29. Example- 1 • The heat of combustion of ethylene at 17° C and at constant volume is -332.19 kcal. Calculate the heat of combustion at constant pressure considering water to be in liquid state (R = 2 cal.). The chemical equation for the combustion of ethylene is C2H4 + 3 O2 = 2CO2(g) + 2H2O (1) 1 mole 3 moles 2moles negligible volume No. of moles of the products = 2 No. of moles of the reactants = 4 Δn =(2-4) =-2 ΔH = ΔE + Δn RT ΔH = -332.19 +[ 2 x I0-3x -2 x 290] =-333.3 kcal Given that ΔE=-332. 19 kcal. T= 273+17= 290k R=2cal=2xlO-3kcals.

30. Example 2 • The heat of combustion of carbon monoxide at constant volume and at 17° C is -283.3 Kj. Calculate its heat of combustion at constant pressure(R= 8.314 J degree-1 mole-1). CO(g) + ½ O2(g) →CO2(g) 1 mole ½ mole 1 mole No. of mles of products = 1 No. of moles of reactants =1.5 n = No. of moles of products - No. of moles of reactants =1-1.5 =-0.5 ΔH = ΔE + Δn x RT ΔH= -283.3 + (-0.5x (8.314x10-3) x 290] = - 283.3-1.20 =-284.5 KJ Heat of combustion of CO at constant pressure is -284.5 kJ. Given that : ΔE =-283.3 kJ T = (273+17) = 290 K. R = 8.314 J or 8.314x10-3 KJ

31. Problem • How much heat could be obtained by the combustion of 10.0 grams CH4burning in the presence of oxygen at constant pressure.

32. Thermochemical equations • factors which affects the quantity of heat evolved or absorbed during a physical or a chemical transformation. • Amount of the reactants and products • Physical state of the reactants and products • Temperature • Pressure

33. Thermochemical equations It must essentially: • be balanced; • give the value of ΔE or ΔH corresponding to the quantities of substances given by the equation; • mention the physical states of the reactants and products . The physical states are represented by the symbols (s), (L), (g) and (aq) for solid, liquid, gas and aqueous states respectively.

34. Example of thermochemical equation H2 + ½ O2 → H2O ΔH = -68.32 Kcal 1 mole of hydrogen reacts with 0.5 mole of oxygen, one mole of water is formed and 68.32 Kcal of heats evolved at constant pressure. But, not specify whether water is in the form of steam or liquid H2 (g) + ½ O2 (g)→ H2O (L) ΔH = -68.32 Kcal H2 (g) + ½ O2 (g)→ H2O (g) ΔH = -57.80 Kcal. Effect of temperature ?????

35. Standard enthalpy change ΔH°. • The heat change at • 298 K and • one atmosphere pressure is called the standard heat change or standard enthalpy change. It is denoted by ΔH°.

36. How do we relate change in temp. to the energy transferred? Heat capacity (J/oC) = heat supplied (J) temperature (oC) Heat Capacity = heat required to raise temp. of an object by 1oC

37. Specific Heat Capacity (Cs) J / oC / g Heat capacity J / oC = = g Mass Specific heat capacity is the quantity of energy required to change the temperature of a 1g sample of something by 1oC

38. The equations representing the variation of heat changes of reaction with temperature are known as Kirchoff's equations.

39. internal energies of the reactants and products. 1. At constant volume, ΔE=E2-E1 Differentiating this equation with respect to temperature at constant volume, we get Variation of heat (or enthalpy) of reaction with temperatureKirchoff's equations.

40. heat capacities heat capacities of the products and reactants But we have already seen that Integrating between temperature T l and T2 , we have Kirchoff's equations. ΔE2- ΔE1 = Δ E2 -ΔE1 = Δ Cv (T2-T1) (3)

41. ΔH=H2-H1 2. At constant pressure Kirchoff's equations. And finely …. ΔH2- ΔH1= ΔCp (T2-T1) (6)

42. Example 3 – 3 The heat of reaction ½ H2 + ½ Cl2 = HCl at 27 C° is -22.1 Kcal. Calculate the heat of reaction at 77°C. The molar heat capacities at constant pressure at 27° C for hydrogen, chlorine and HCl are 6.82, 7.70 and 6.80 cal mol-1, respectively. Here ½ H2 + ½Cl- → HCl Δ H = -22.1 Kcal Δ Cp = Heat capacities of products - Heat capacities of reactants = 6.80-[½(6.82) +½ (7.70)] = 6.80 - 7.26 = -0.46 x 10-3 Kcal ΔH2- ΔH1= ΔCp (T2-T1) ΔH2 - (-22.1) = (-0.46 x 10 -3) x 50= -21.123 T2 = 273 + 77 =350 K T1 = 273 + 27 = 300 K T2-T1= (350-300)K=50K

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

44. Specific Heat Capacity How much energy is transferred due to T difference? The heat (q) “lost” or “gained” is related to a) sample mass b) change in T and c) specific heat capacity

45. Specific Heat Capacity Substance Spec. Heat (J/g•K) H2O 4.184 Ethylene glycol 2.39 Al 0.897 glass 0.84 Aluminum

46. 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?

47. heat gain/lose = q = (sp. ht.)(mass)(∆T) 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? where ∆T = Tfinal - Tinitial q = (0.897 J/g•K)(25.0 g)(37 - 310)K q = - 6120 J Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

48. We define the standard state of a substance as the state of the pure substance at 1 atm pressure and the temperature of interest (usually 25 °C). • The standard enthalpy change (ΔH°) for a reaction is the enthalpy change in which reactants and products are in their standard states. • The standard enthalpy of formation (ΔHf°) for a reaction is the enthalpy change that occurs when 1 mol of a substance is formed from its component elements in their standard states.

50. End of par A