1. AP Chemistry 12 Energy Relationships in Chemistry

2. Thermochemistry Thermodynamics – the study of energy and its transformations. Thermochemical changes – energy changes associated with chemical reactions. System  that specify part of the universe of interest to us. Surroundings  the part of the universe not contained in the system.

3. 3 types of Systems open system  exchanges mass and energy closed system  exchanges energy but no mass isolated system  no exchange of either mass or energy

4. Three Types of Systems Open system Closed System cork Isolated System insulation

5. Different Types of Energy Energy – the ability to do work. Thermal energy – associated with the random motions of atoms and molecules Heat energy – transfer of thermal energy between two objects at different temperature.

6. Energy (cont’d) Chemical energy – energy stored within the structural units of chemical substance. Potential energy – the ability of an object to do work because of its position in a field of force.

7. Kinetic Energy – the work that can be performed by a moving object. The unit of energy 1 Joule (J) =1 kg m 2 /s 2 An older unit of energy 1 calorie (cal) = 4.184 J exactly

8. The Law of Conservation of Energy The law of conservation of energy  Energy is neither created nor destroyed in ordinary chemical and physical processes.  Converted from one type into another.

9. This is also stated in terms of the first law of thermodynamics.  E = internal energy change of the system E f and E i  the energy of the final and initial states, respectively

10. First Law of Thermodynamics Chemical reactions either absorb or release energy. Two terms  Exothermic reaction  heat is released to the surroundings.  Endothermic reaction  heat is supplied to the system by the surroundings.

11. Exothermic

12. Endothermic

13. The First Law Restated chemical systems – examine the conversion of heat energy into work.

14. Signs for Heat and Work Work done by system on surroundings  w ‘-’ Work done by surroundings on system  w ‘+’ q < 0, heat flows to surroundings  Exothermic ‘-’ q > 0, heat flows to system  Endothermic ‘+’

15. State and Path Functions  E,  H,  V are examples of state functions.  State functions – numerical value doesn’t depend on how the process is carried out. Work (w) and q (heat) are path functions  The amount of work done or heat released depends on how the system changes states.

16. Enthalpies of Formation – Standard Reaction Enthalpies The enthalpy change for the reaction  H rxn =  H(products) -  H(reactants) We cannot measure the absolute values of the enthalpies!! How do we ‘measure’ enthalpies (or heat contents) of chemical species?

17. The Formation Reaction A "chemical thermodynamic reference point." For CO and CO 2 C (s) + O 2 (g)  CO 2 (g) C (s) + ½ O 2 (g)  CO (g) The "formation" of CO and CO 2 from its constituent elements in their standard states under standard conditions.

18. The Formation Reaction The formation reaction For the formation of 1.00 mole of Na 2 SO 3 (s) 2 Na(s) + S(s) + 3/2 O 2 (g)  Na 2 SO 3 (s) The ‘formation enthalpy of Na 2 SO 3 (s)’, symbolised  H f  [Na 2 SO 3 (s)]

19. Standard Conditions for Thermodynamic Reactions The degree sign, either  or , indicates standard conditions  P = 1.00 atm  [aqueous species] = 1.00 mol/L  T = temperature of interest (note 25  C or 298.15 K is used in the tables in your text).

20. The Significance of the Formation Enthalpy  H f ° is a measurable quantity! Compare CO (g) with CO 2 (g) C (s) + 1/2 O 2 (g)  CO (g)  H f ° [CO(g)] = -110.5 kJ/mole C (s) + O 2 (g)  CO 2 (g)  H f ° [CO 2 (g)] = - 393.5 kJ/mole The formation enthalpy for CO 2 (g) is larger than the formation enthalpy of CO (g).

21. Reaction Enthalpies Formation enthalpies – thermodynamic reference point, Formation of the elements from themselves is a null reaction –  H f  (elements) = 0 kJ / mole.

22. The Combustion of Propane

23. The General Equation Calculate enthalpy changes from the formation enthalpies as follows. Reverse a reaction, the sign of the enthalpy change for the reaction is reversed. Multiply a reaction by an integer, the enthalpy change is multiplied by the same integer.

24. The Measurement of Energy Changes – Calorimetry Calorimetry – the measurement of heat and energy changes in chemical and physical processes. Heat capacity (C) – the amount of heat (energy) needed to raise the temperature of a given mass of substance by 1°C. Specific heat capacity (s) – the amount of heat energy (in Joules, J) required to raise 1 g of a substance by 1°C (units = J/g °C).

25. General expression for heat capacity C = m s  m is the mass of the substance (in grams). Molar heat capacity C m = M s  M – molar mass of the substance  s – its specific heat capacity.

26. The Calorimeter A calorimeter – a device which contains water and/or another substance with a known capacity for absorbing energy (heat). Calorimeters are adiabatic systems. All energy changes take place within the calorimeter.

27. Adiabatic System Adiabatic system – thermally insulated from the rest of the universe No heat exchange between system and surroundings! For an adiabatic system, q total = q rxn + q H 2 O + q cal = 0  -q rxn = q H 2 O + q cal

28. The Constant Volume (Bomb) Calorimeter  E = q v

29. The Constant Pressure Calorimeter  H = q p

30. Other important Enthalpy changes Many other important processes have associated enthalpy changes. The measurement of the heat changes for these process can give us some insight into the changes in intermolecular forces that occur during the transformation.

31. Heat of dilution and solution.  H sol = the heat absorbed or given off when a quantity of solute is dissolved in a solvent.  H sol = H(sol’n) - H(component)  H(component) = H (solid) + H(solvent)

32. For the process, HCl (aq, 6 M)  HCl (aq, 1 M). A significant amount of heat is released when the acid solution is diluted. This is the enthalpy of dilution of the acid.  H dil = H(sol’n 2) – H(sol’n,1)

33. Lattice Enthalpies Look at the following process. NaCl (s)  Na + (g) + Cl - (g)  H =  H lat = 788 kJ/mole  the lattice enthalpy A very endothermic reaction! Due to the strength of the ionic bond!

34. Latent Heats Latent heats are the enthalpy changes associated with phase transitions. H 2 O (l)  H 2 O (g)  H r =  H vap  the enthalpy of vapourization. H 2 O (s)  H 2 O (l)  H r =  H fus  the enthalpy of fusion. H 2 O (s)  H 2 O (g)  H r =  H sub  the enthalpy of sublimation.

35. Latent Heats

36. Foods and Fuels Most of the chemical reactions that produce heat are combustion reactions. Note – all combustion reactions are exothermic. Fuel values are generally reported as positive quantities. Obtaining fuel values – calorimetry.

37. Fossil Fuels Coal, petroleum, and natural gas are known as fossil fuels. They are collectively the major source of energy for commercial and personal consumption. Fossil fuels are mixtures of many different kinds of organic compounds. The fuel values of fossil fuels is directly related to the amount of carbon and hydrogen in the fuel.

38. Hydrogen As a Fuel Hydrogen has a huge fuel value (142 kJ/g). The combustion product is innocuous – water. Obviously, there are problems! Two major difficulties with H 2 as a fuel source.  Where do we get the hydrogen?  How do we store the hydrogen?