Thermochemistry

1. Energy in Chemical Reactions Thermochemistry

2. Name a chemical reaction that releases heat or light energy. Which would have more energy in those cases, the reactants or products? Think of reaction that absorbs heat and think about the reactants and products energy. Comparisons of Energy release http://www.artisanbreadbaking.com/images/altamura/altamura_13.jpg http://images.search.yahoo.com/images/view?back=http%3A%2F%2Fimages.search.yahoo.com%2Fsearch%2Fimages%3Fp%3Dcandle%2Bburning%26

3. “Thermo” means “heat” “Dynamics” means “power: The study of energy and its transformations Energy- the capacity to do work or to transfer heat Work-the energy used to cause an object with mass to move against a force Heat- energy used to cause the temperature an object to increase Thermodynamics

4. Kinetic –energy of motion • Ek = ½ mv2 m = mass v = velocity • Potential – energy of position • Ep = mgh m=mass g=gravity(9.8m/s2) h=height • Electrostatic energy- opposite charges attract each other, and like charges repel each other • Eel =κ Q1 Q2κ = proportionality constant 8.99x109J-m/C2 • D Q=electrical charges magnitude of electron D = distance Energy (p 8-9)

5. Example: Electrostatic

6. System- the part of the universe of interest Surroundings-everything relevant to system Example: system might be reactants in a beaker while the surroundings would be the room temperature, pressure and humidity Energy Flow

7. Energy within the system is found by: ∆E =Efinal – Einitial = Eproducts - Ereactants ∆E = ∆U : internal energy If Eproducts > Ereactants E= “+” If Eproducts < EreactantsE= “-” Internal Energy

8. Heat – q; Energy transferred between a system and it’s surroundings due to ______difference Work- w; Energy transferred when an object is moved by a force ∆E = q + w Energy moving into a system is “ “ Energy moving out of a system is “ “ Heat – Thermal Energy/Work + -

9. Sum of internal energy plus the product of the pressure and volume ∆H = ∆E + P∆V (remember liquids and solids undergo negligible volume change) (note that if gases are present-volume ∆ small!) Therefore- often ∆H = ∆E or close to it Enthalpy

10. Heat change absorbed by water can be measured as: Specific Heat Capacity (c)- Which had higher c : Al vs. water ? Calorimetry q = c x mass x ∆T q = quantity of heat c = specific heat capacity ∆T = change in temperature

11. Heats of reaction correspond in magnitude to the amount reacted in moles and their states of matter (solid/liquid/gas) For example: 2H2O(l)  2H2(g) + O2(g) ∆Hrxn= 572kJ Reverse: 2H2(g) +O2(g)  2H2O(l) =-572kJ Or… H2(g) + ½ O2(g)  H2O(l) ∆Hrxn=-286kJ Stoichiometry in Thermochem Do all our reactions with exactly one mole? Create a potential problem with one of the above 3 equations giving the mass which would then be converted into moles. How much energy would be absorbed or released?

12. The overall enthalpy is the sum of enthalpy steps leading to the final product. The enthalpy of the combustion of C to CO2 is -393.5 kJ/mol C and the enthalpy for the combustion of CO2 is -283.0 kJ/mol CO: (1) C(s) + O2(g)  CO2(g) ∆H1 = -393.5 kJ (2) CO(g) + ½ O2  CO2(g) ∆H2 = -283.0 kJ This data can be found in Appendix B Calculate the enthalpy for the combustion of C to CO: (3) C(s) + ½ O2(g)  CO(g) ∆H3 = Hess’s Law of Summation -110.5 kJ

13. Gases- are at 1 atm pressure • Aqueous solutions-are 1 M concentration • Pure substances-1 atm pressure; 298 K— • What state are they in now? • Standard Heat of formation ∆Hf • Take apart the equation pg 254…what is the enthalpy? Standard State Assumptions:

14. ∆Hrxn= ∑ m∆Hf - ∑ n∆Hi Try to find the total enthalpy of the reaction: 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(l) Sum Total!

15. Entropy: S ∆Srxn = ∑ m Sproducts - ∑ n Sreactants How are they similar? Look again… Calculate ∆S for the synthesis of ammonia from N2(g) and H2(g) at 298 K N2(g) + 3 H2(g)  2 NH3(g) ∆S = -198.3 J/K

16. Temperature increases causes entropy increase –directly proportional • Physical state/phase change-as a compd absorbs heat (q>0) S˚increases • Dissolving a solid/liquid-increases disorder and therefore S˚ • Dissolving a gas dispersion increases S˚ • Increase in atomic size/complexity increases Predicting Entropy Signs

17. In Exothermic- (∆H < 0) the heat released by the system increases the total entropy (S) of the system + surroundings: ∆Suniverse > 0; ∆Ssystem + ∆Ssurroundings > 0 • In Exothermic- (∆H < 0) the heat released by the system increases the entropy of the surroundings such that it overwhelms the decrease in the system then ∆Suniverse>0 Spontaneous Exothermic ReactioNS

18. For an endothermic reaction the ∆Suniverse is positive if the ∆Ssystem large enough (∆S >>0) to overwhelm the ∆Ssurroundings (∆Ssurroundings<0) Spontaneous Endothermic reactions http://patrickking.org/tug.htm

19. Brainstorm an equation that would combine the concepts of entropy & enthalpy to determine if a reaction is spontaneous- Spontaneous value? Enter Gibbs Free Energy! ∆G = measure of the spontaneity of a process and of the useful energy available from it. G = H – TS H = Enthalpy T= Temperature S = Entropy THE SIGN OF G TELLS US WHETHER A REACTION IS SPONTANEOUS!

20. ∆Suniverse>0 Spontaneous • ∆Suniverse<0 Non-spontaneous • ∆Suniverse =0 Equilibrium • ∆G < 0 Spontaneous • ∆G > 0 Non-spontaneous • ∆G = 0 Equilbrium Signs of S & G

21. ∆G˚ = ∆H˚sys - T∆S˚sys • Also: • ∆G˚ = ∑m∆G˚f(products) - ∑n∆G˚f(reactants) Calculating Gibbs Free Energy With your shoulder partner, read Sample Problem 20.4. After attempting it and following through the solution. Try Follow-up Problem 20.4

22. http://www.youtube.com/watch?v=OsldTzBaIEw http://www.youtube.com/watch?v=V48r4IArzLc&feature=related Entropy/Gibbs Fun!