Thermochemistry

1. Thermochemistry

2. Energyis the capacity to do work. • Radiant energy comes from the sun and is earth’s primary energy source • 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 • Potential energy is the energy available by virtue of an object’s position • Kinetic energy is the energy of motion

3. Temperature = Thermal Energy 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.

4. 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. • Surrounding: • the rest of the universe. open isolated closed energy nothing Exchange: mass & energy

5. Collision Theory • To explain factors affecting rxn rate • For two molecules to react, they must come in contact or collide. • When molecules collide, • Ineffective collision occurs: molecules may not orient in right position. No Rxn. • Effective collision occurs: molecules possess enough kinetic energy to position.

6. What Happen Just Before Rxn • Molecules approaches each other. • Repulsive force increases as they approach each other. • There is a minimum energy to push the molecules to collide Activation Energy (Ea) • Supplying enough kinetic energy will “kick-start” the reaction.

7. Graphing Rxn: Spontaneous

8. Gibbs Free Energy (ΔG) • If ΔG is positive, or ΔG > 0Indication: Rxn  Nonspontaneous • If ΔG is negative, or ΔG < 0Indication: System  Spontaneous • The previous example is spontaneous.

9. Rate Influencing Factor • Surface Area • Temperature • Concentration • Presence of Catalysts

10. Factor 1: Surface Area • Similar to solubility • More surface area enhances collision frequency • Homogeneous Rxn • The reactants are in the same physical state.For example, gas reacting with gas • Heterogeneous Rxn • The reactants are in DIFFERENT physical stateFor example: Solid reacting with liquid/aq. soln

11. Factor 2: Temperature • Increase in temperature provide more kinetic energy for reaction to proceed • Fast-moving molecules have better chance of “effective collision.” • It also enhances frequency of collision. • Refrigerator is designed to SLOW reaction rate.

12. Factor 3: Concentration • Higher chemical concentration generally promotes reaction rate • Concentration is directly proportional to the frequency of collision, in general. • Just like medicine, we need higher dosage for severely-ill patient.

13. Factor 4: Catalyst • In human body, we have enzymes acting as catalyst. • It does not undergo permanent change. • It provides an alternative reaction pathway of Lower Activation Energy. • Inhibitor: reduces a reaction rate by preventing the reaction occurring in the usual way. • Example: Food preservatives

14. Effect of Catalyst on Rxn

15. The First Law of Thermodynamics • Relating DE to Heat(q) and Work(w) • Energy cannot be created or destroyed. • Energy of (system + surroundings) is constant. • Any energy transferred from a system must be transferred to the surroundings (and vice versa). • From the first law of thermodynamics:

16. First Law of Thermodynamics Energy can be converted from one form to another but energy cannot be created or destroyed. Second Law of Thermodynamics The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process. Spontaneous process: DSuniv = DSsys + DSsurr > 0 Equilibrium process: DSuniv = DSsys + DSsurr = 0

17. Enthalpy Changes The difference between the potential energy of the reactants and the products during a physical or chemical change is the Enthalpy change or ∆H. AKA: Heat of Reaction at constant pressure

18. H2O (s) H2O (l) DH = 6.01 kJ Thermochemical Equations Is DH negative or positive? System absorbs heat Endothermic DH > 0 6.01 kJ are absorbed for every 1 mole of ice that melts at 00C and 1 atm.

19. DH = -890.4 kJ CH4(g) + 2O2(g) CO2(g) + 2H2O (l) Thermochemical Equations Is DH negative or positive? System gives off heat Exothermic DH < 0 890.4 kJ are released for every 1 mole of methane that is combusted at 250C and 1 atm.

20. Graphing Rxn: Endothermic The temperature goes down

21. Graphing Rxn: Exothermic The temperature goes up

22. Cold and hot packs • How do instant hot and cold packs work?

23. Hot pack • Pressing the bottom , the diaphragm breaks. • Calcium chloride dissolves in water and warms it. • The beverage gets warm.

24. Exothermic process • Heat flows into the surroundings from the system in an exothermic process. Energy Surroundings Hot pack Temperature rises

25. Cold pack • Water and ammonium nitrate are kept in separate compartments. • Pressing the wrapper, the ammonium nitrate dissolves in water and absorbs heat. • The pack becomes cold. It is used to treat sports injuries.

26. Endothermic process • Heat flows into the system from the surroundings in an endothermic process. Surroundings Cold pack Energy Temperature falls

27. Explosions This reaction is exothermic!

28. Photosyntesis This reaction is endothermic!

29. Combustions These reactions are exothermic!

30. Enthalpy Changes in Reactions All chemical reactions require bond breaking in reactants followed by bond making to form products Bond breaking requires energy (endothermic) while bond formation releases energy (exothermic)

32. Enthalpy Changes in Reactions endothermic reaction - the energy required to break bonds is greater than the energy released when bonds form. exothermic reaction - the energy required to break bonds is less than the energy released when bonds form.

33. Enthalpy Changes in Reactions ∆H can represent the enthalpy change for a number of processes Chemical reactions ∆Hrxn – enthalpy of reaction ∆Hcomb – enthalpy of combustion

34. Formation of compounds from elements ∆Hof– standard enthalpy of formation The standard molar enthalpy of formation is the energy released or absorbed when one mole of a compound is formed directly from the elements in their standard states.

35. Phase Changes ∆Hvap – enthalpy of vaporization ∆Hfus – enthalpy of melting ∆Hcond – enthalpy of condensation ∆Hfre – enthalpy of freezing Solution Formation ∆Hsoln – enthalpy of solution

36. q = mc∆T 140 100 q = n∆H Temp. (°C ) q = mc∆T q = n∆H 0 q = mc∆T -40 Time

37. Data: cice = 2.01 J/g.°C cwater = 4.184 J/g.°C csteam = 2.01 J/g.°C ΔHfus = +6.02 kJ/mol ΔHvap= +40.7 kJ/mol

38. warming ice: q = mc∆T = (40.0)(2.01)(0 - -40) = 3216 J warming water: q = mc∆T = (40.0)(4.184)(100 – 0) = 16736 J warming steam: q = mc∆T = (40.0)(2.01)(140 -100) = 3216 J

39. melting ice: q = n∆H = (2.22 mol)(6.02 kJ/mol) = 13.364 kJ boiling water: q = n∆H = (2.22 mol)(40.7 kJ/mol) = 90.354 kJ n = 40.0 g 18.02 g/mol = 2.22 mol moles of water:

40. Total Energy 90.354 kJ 13.364 kJ 3216 J 3216 J 16736 J 127 kJ

41. Sample Problem • How much heat is required to change 36 g of H2O from -8 deg C to 120 deg C? Step 1: Heat the ice Q=mcΔT Q = 36 g x 2.06 J/g deg C x 8 deg C = 593.28 J = 0.59 kJ Step 2: Convert the solid to liquid ΔH fusion Q = 2.0 mol x 6.01 kJ/mol = 12 kJ Step 3: Heat the liquid Q=mcΔT Q = 36g x 4.184 J/g deg C x 100 deg C = 15063 J = 15 kJ

42. Sample Problem • How much heat is required to change 36 g of H2O from -8 deg C to 120 deg C? Step 4: Convert the liquid to gas ΔH vaporization Q = 2.0 mol x 44.01 kJ/mol = 88 kJ Step 5: Heat the gas Q=mcΔT Q = 36 g x 2.02 J/g deg C x 20 deg C = 1454.4 J = 1.5 kJ Now, add all the steps together 0.59 kJ + 12 kJ + 15 kJ + 88 kJ + 1.5 kJ = 118 kJ

43. specific heat capacity - the amount of energy , in Joules (J,SI unit), needed to change the temperature of one gram (g) of a substance by one degree Celsius (°C). c: does not depend on mass, so intensive property. The symbol for specific heat capacity is a lowercase c The unit is J/ g°C or kJ/ g°C Heat/Enthalpy Calculations

44. A substance with a large value of c can absorb or release more energy than a substance with a small value of c. with the larger c will undergo a smaller temperature change with the same amount of heat applied.

45. FORMULA q = mc∆T q = heat (J) m = mass (g) c = specific heat capacity ∆T = temperature change = T2 – T1 = Tf – Ti

46. eg. How much heat is needed to raise the temperature of 500.0 g of water from 20.0 °C to 45.0 °C? Solve q = m c ∆T for c, m, ∆T, T2 & T1

47. heat capacity- the quantity of energy , in Joules (J, SI unit), needed to change the temperature of a substance by one degree Celsius (°C) The symbol for heat capacity is uppercase C C: depends on mass, so extensive property. The unit is J/ °C or kJ/ °C

48. FORMULA C = mc q = C ∆T C = heat capacity c = specific heat capacity m = mass ∆T = T2 – T1

49. Calorimetry • Heat Capacity and Specific Heat • Calorimetry = measurement of heat flow (follow Law of conservation of energy). • Heat lost by one substance = Heat gained by other substance • Calorimeter = apparatus that measures heat flow. • Heat capacity = the amount of energy required to raise the temperature of an object (by one degree). • Molar heat capacity = heat capacity of 1 mol of a substance. • Specific heat = specific heat capacity = heat capacity of 1 g of a substance.

50. Calorimetry Heat (q) absorbed or released: q = msDt The heat capacity (C) of a substance q = CDt C = ms Dt = tfinal - tinitial Constant Pressure Calorimetry