Thermochemistry : The Study of heat change in chemical reactions

1. Topic 6: Part I Thermochemistry Thermochemistry: The Study of heat change in chemical reactions

2. Energy – the capacity to do work or produce heat • Types of Energy: • Radiant – solar energy, comes from the sun • Thermal – energy associated with the random motion of atoms and molecules • Chemical – energy stored within the structural units of chemical substances • Potential – energy available by virtue of an object’s position • Every chemical reaction obeys two fundamental Laws: (1) The Law of Conservation of Mass, and (2) The Law of Conservation of Energy

3. The nature of energy • Energy is the capacity to do work or produce heat. • Law of Conservation of Energy: energy can be converted from one form to another, but not created or destroyed. • Potential – energy of position or composition • Kinetic – energy of motion • Temperature reflects the avg KE of the molecules. • Heat involves the transfer of energy between two objects.

4. Energy is …… • Change in energy (∆E) is calculated using heat (q) and work (w). • It is a state function – meaning it is independent of the pathway, or how you get from point A to B. • Work is a force acting over a distance. • Heat is energy transferred between objects because of temperature difference.

5. Chemical Energy in Reactions ( keeping track ) • The Universe is divided into two parts: • the system (sometimes a reaction) • and the surroundings. • The system is the part we are concerned with. • The surroundings are the rest. • Exothermic reactions release energy to the surroundings. • Endothermic reactions absorb energy from the surroundings.

6. Heat Potential energy

7. Heat Potential energy

8. Direction • Every energy measurement has three parts. • A unit ( Joule, J= kg∙m2 / s2 ). • A number indicating how many • A sign to tell direction. • negative - exothermic • positive - endothermic

9. Surroundings System Energy DE <0

10. Surroundings System Energy DE >0

11. We apply the same rules of energy change for heat and work. • Heat given off is negative. • Heat absorbed is positive. • Work done by system on surroundings is negative. • Work done on system by surroundings is positive. • Thermodynamics- The study of energy and the changes it undergoes.

12. First Law of Thermodynamics • The energy of the universe is constant. • Law of conservation of energy. • q = heat • w = work • DE = q + w • Take the system’s point of view to decide signs.

13. Enthalpy • abbreviated H • H = E + PV (that’s the definition) • E = internal energy of a system • P = pressure of the system • V = volume of the system • at constant pressure. • DH = DE + PDV

14. Enthalpy Con’t • the heat, q, at constant pressure can be calculated from • DE = q + w (w = - PDV) • Substitute for work (w) and solve for heat (q) • q = DE + P DV = DH • Therefore q = DH at constant pressure • We use DH as a measure of the change in enthalpy or heat in joules of a system.

15. Enthalpy con’t • Heat is a measure of change in enthalpy of a system at constant P. • Heat of a reaction = change in enthalpy • For a chemical reaction: • ∆H = Hproducts – Hreactants • + ∆H = heat absorbed by system (endothermic) • - ∆H = heat lost by system (exothermic)

16. Example • 6.4 When one mole of methane (CH4) is burned at constant pressure, 890 kJ of energy is released as heat. Calculate ∆H for a process in which a 5.8g sample of methane is burned at constant pressure.

17. Calorimetry • Measuring heat. • Use a calorimeter. • Two kinds • Constant pressure calorimeter (called a coffee cup calorimeter) • heat capacity for a material, C, is the measure of the energy needed to raise the temp. of an object 1oC. • C= heat absorbed/ DT = DH/ DT

18. Calorimetry • specific heat capacity, s, is the heat capacity given per gram of substance (C/g). • molar heat capacity = C/moles • the heat of reaction, q = s x m x DT • S = specific heat capacity • m = mass of substance • DT = change in temperature

19. Calorimetry • A coffee cup calorimeter measures DH. • An insulated cup, full of water. • The specific heat of water is 4.18J/ºC∙g • Heat of reaction= DH = s x m x DT

20. Examples • If 50ml of 1.0M HCl at 25oC is reacted with 50ml of 1.0M NaOH at 25oC. The final temperatrue of this exothermic reaction is 31.9oC. • Calculate the change in heat • Calculate ∆H in kJ/mol

21. Example • When 1.00 L of 1.00 M Ba(NO3)2 solution at 25oC is mixed with 1.00L of 1.00M Na2SO4 solution at 25oC in a calorimeter, the white solid BaSO4 forms and the temperature of the mixture increases to 28.1oC. • Assuming the calorimeter absorbs only a negligible quantity of heat, that the specific heat capacity of the solution is 4.18 J/oC ∙ g and that the density of the final solution is 1.0g/mL, calculate the enthalpy change per mole of BaSO4 formed.

22. Calorimetry • Constant volume calorimeter is called a bomb calorimeter. • Material is put in a container with pure oxygen. Wires are used to start the combustion. The container is put into a container of water. • The heat capacity of the calorimeter is known and tested. • Since DV = 0, PDV = 0, DE = q

23. Bomb Calorimeter • thermometer • stirrer • full of water • ignition wire • Steel bomb • sample

24. Properties • intensive properties not related to the amount of substance. • density, specific heat, temperature. • Extensive property - does depend on the amount of stuff. • Heat capacity, mass, heat from a reaction.

25. Hess’s Law • Enthalpy is a state function meaning it is independent of the path. • We can add equations to come up with the desired final reaction and add the associated DH values to get the final ∆H. • Two rules using Hess’s Law: • If the reaction is reversed the sign of DH is changed • If the reaction is multiplied by a constant, so is DH for that reaction.

26. Nitrogen + Oxygen → Nitrogen Dioxide Overall Reaction: N2 + 2O2 → 2NO2∆H1 = 68kJ Reaction steps added: N2 + O2 → 2NO ∆H2 = 180kJ 2NO + O2 → 2NO2 ∆H3 = -112kJ ____________________________________________________________________________________________________________________________________________________ N2 + 2O2 → 2NO2∆H2 + ∆H3 = 68kJ

27. O2 NO2 -112 kJ H (kJ) 180 kJ NO2 68 kJ N2 2O2

28. Characteristics of Enthalpy Changes • If a reaction is reversed the sign of ∆H is also reversed. • The magnitude of ∆H is directly proportional to the quantities of reactants and products. • If coefficients are multiplied by an integer, then ∆H is multiplied by the same integer • Examples: 6.7 and 6.8

29. Standard Enthalpy • The enthalpy change for a reaction at standard conditions (25ºC, 1 atm , 1 M solutions) • Symbol DHº • When using Hess’s Law, work by adding the equations up to make it look like the answer. • The other parts will cancel out.

30. Standard Enthalpies (Heat) of Formation • Hess’s Law is much more useful if you know lots of reactions. • Made a table of standard heats of formation. The amount of heat needed to form 1 mole of a compound from its elements in their standard states. • Standard states are 1 atm, 1M and 25ºC • For an element it is 0 • There is a table in Appendix 4 (pg A21)

31. Standard Enthalpies (Heat) of Formation • Need to be able to write the equations. • What is the equation for the formation of NO2 ? • ½N2 (g) + O2 (g) ® NO2 (g) • Have to make one mole to meet the definition. • Write the equation for the formation of methanol CH3OH.

32. Since we can manipulate the equations • We can use heats of formation to figure out the heat of reaction.

33. Examples 6.9 and 6.10 • Using the standard enthalpies of formation listed in the appendix, calculate the standard enthalpy change for the overall reaction that occurs when ammonia is burned in air to form nitrogen dioxide and water. • 4NH3(g) + 7O2(g) → 4NO2(g) + 6H2O (l) • Using enthalpies of formation, calculate the standard change in enthalpy for the thermite reaction: • 2Al (s) + Fe2O3(s) → Al2O3(s) + 2Fe (s)