1. Energy and Chemical Reactions Energy is transferred during chemical and physical changes, most commonly in the form of heat

2. Energy Energy can be kinetic – associated with motion, such as thermal, mechanical, electric, sound Energy can be potential – associated with an object ’ s position, such as chemical, gravitational, electrostatic Energy is converted from one form to another

3. First Law of Thermodynamics The total energy of the universe is constant Energy is conserved

4. Temperature and Heat Temperature is a measure of the average kinetic energies of the particles in a substance Heat is energy that can be transferred between substances that are at different temperatures Heat will transfer between two objects in contact until thermal equilibrium occurs

5. Heat transfer The quantity of heat lost by a hotter object and the quantity of heat gained by a cooler object when they are in contact are numerically equal (but opposite direction) Exothermic – heat is transferred from the system to the surroundings Endothermic – heat is transferred from the surroundings to the system

6. Energy Units Joule is the SI unit for thermal energy 1 J = 1 kg. m 2 /s 2 Kilojoules are also commonly used The calorie is an older unit for heat; 1 cal = 4.184 J Dietary Calories are actually 1000 calories

7. Specific Heat Capacity and Heat Transfer The quantity of heat transferred to or from an object when its temperature changes depends on: –Quantity of the material –Size of the temperature change –Identity of the material Specific heat capacity – the quantity of heat required to raise the temperature of 1.00g of a substance by one kelvin (J/g. K)

8. q = m c  T q is heat in joules m is mass in grams T = T final – T initial Water has a particularly high specific heat; metals have low specific heats

9. Assumptions Heat transfers until both substances are at the same temperature We assume no heat is transferred to warm the surroundings (though this is not accurate) The heat that is lost by one substance is equal and opposite in sign to the heat that is gained by the other substance

10. Energy and Changes of State Heat of fusion – energy to convert a substance from solid to liquid (J/g) Heat of vaporization – energy to convert a substance from liquid to gas (J/g) The energy required for a change of state is determined by the type of substance and its quantity (mass)

12. Calorimetry Constant pressure calorimetry measures H Constant pressure calorimetry can be done with a coffee-cup calorimeter A reaction changes the temperature of the solution in the calorimeter; measuring the change in the solution allows calculation of the change in the reaction q rxn + q solution = 0

13. Calorimetry Constant volume calorimetry measures E A bomb calorimeter is used for constant volume calorimetry q rxn +q bomb +q water = 0

14. Thermodynamics – the study of heat and work  E = q + w  E is the change in kinetic and potential energies of the system Positive q is heat going into the system Negative q is heat leaving the system Positive w is work done on the system Negative w is work done by the system Work (of a gas): w = - P(  V)

15. State Functions A quantity that is the same no matter what path is chosen in going from initial to final Changes in internal energy and enthalpy for chemical or physical changes are state functions Neither heat nor work individually are state functions, but their sum is

16. Enthalpy Changes for Chemical Reactions Measures the change in heat content Enthalpy changes are specific to the identity and states of reactants and products and their amounts H is negative for exothermic reactions and positive for endothermic reactions Values of H are numerically equal but opposite in sign for chemical reactions that are the reverse of each other Enthalpy change depends on molar amounts of reactants and products

17. 2 Methods to find  H rxn Hess’s Law (indirect method) If a reaction is the sum of two or more other reactions, H for the overall process is the sum of the H values of those reactions

18. C(s) + O 2 (g)  CO 2 (g) H=-393.5 kJ/mol CO(g) + 1/2O 2 (g)  CO 2 (g) H=-283.0 kJ/mol _________________________________ C(s) + 1/2O 2 (g)  CO(g) H=?

19. C(s) + O 2 (g)  CO 2 (g) H=-393.5 kJ/mol H 2 (g) + 1/2O 2 (g)  H 2 O(l) H=-285.8 kJ/mol 2C 2 H 2 (g) + 5O 2 (g)  4CO 2 (g) + 2H 2 O(l ) H=-2598.8 kJ/mol __________________________________ 2C(s) + H 2 (g)  C 2 H 2 (g) H=?

20. Standard Enthalpies of Formation The standard molar enthalpy of formation (H f o ) is the enthalpy change for the formations of 1 mol of a compound directly from its component elements in their standard states The standard state of an element or a compound is the most stable form of the substance in the physical state that exists at standard atmosphere at a specified temperature

21. Standard Enthalpy of Formation The standard enthalpy of formation for an element in its standard state is zero Most enthalpies of formation values are negative, indicating an exothermic process The most stable compounds have the largest exothermic values

22. Enthalpies of Formation Enthalpy change for a reaction can be calculated from the enthalpies of formation of the products and reactants (direct method):  f o (products)] –  f o (reactants)] =  rxn o Reactions with negative values of  rxn o are generally product-favored, while positive  rxn o usually indicates a reactant-favored reaction

23. 2H 2 S(g) + 3O 2 (g)  2H 2 O(l) + 2SO 2 (g) H=? SubstanceH f (kJ/mol) H 2 S (g)-20.15 H 2 O (l)-285.8 SO 2 (g)-296.1

24. Application of Enthalpy 2 Al(s) + Fe 2 O 3 (s)  Al 2 O 3 (s) + 2 Fe(l) 2 Al(s) + Fe 2 O 3 (s)  Al 2 O 3 (s) + 2 Fe(l) H rxn = -822.8 kJ How much heat is released if 1 mole of Al is used? How much heat is released if 1 mole of Al is used? How much heat is released if 4.2 moles of Al is used? How much heat is released if 4.2 moles of Al is used? How much heat is released if 150. g of Al is used? How much heat is released if 150. g of Al is used?