1. Chapter 6 Thermochemistry

2. Energy is... – the capacity to do work or produce heat – conserved, amount of energy in universe never changes – Temperature is measure of how hot or cold an object is – Heat ALWAYS move from hot to cold – State Function or Property - something that is independent of the path, or how you get from point A to B – Work * (w) = force acting over a distance – Heat * (q) = energy transferred between objects * Not a State Function

3. State Function or Property All gases are at 1 atm All liquids are pure All solids are pure All solutions are at 1 M concentrations The energy of formation of an element in its normal state is defined as zero The temperature used for standard state values is almost invariably room temperature: 25 o C (298 K)

4. Kinetic Energy Energy an object possesses by virtue of its motion.

5. Potential Energy Energy an object possesses by virtue of its position or chemical composition.

6. The Universe for Chemists Divided into two halves: 1.System 2.Surroundings The System is the part you are concerned with The Surroundings are the rest Exothermic reactions release energy to the surroundings Endothermic reactions absorb energy from the surroundings

7. System and Surroundings The System includes the molecules we want to study (here gases in system) The Surroundings are everything else (here, the cylinder and piston).

8. Internal Energy (E) Heat Exothermic Chemical Energy

9. Internal Energy (E) Heat Endothermic Chemical Energy

10. Energy Measurement Every energy measurement has three parts: 1. A unit ( Joules or calories). 2. A number, it’s magnitude. 3. A sign to tell direction. Exothermic measurements are negative Endothermic measurements are positive Sign reflects system point of view

11. Units of Energy The SI unit of energy is the joule (J). An older, non-SI unit is still in widespread use: calorie (cal) 1 cal = 4.184 J

12. First Law of Thermodynamics Energy of the universe is constant. Energy can neither be created nor destroyed * Law of conservation of energy. Δ E = q + w w = work q = heat Δ E = KE + PE = Change in Energy * You can’t win – all you can do is break even

13. Internal Energy The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E. Internal energy can be changed by the flow of work, heat or both. ∆ means change in the systems internal energy

14. Sign Conventions for q, w, and ∆E + (plus)- (minus) q sys. gains heatsys. loses heat w work done on sys.work done by sys. ∆E net gain of energy by sys. net loss of energy by sys.

15. Exothermic Reaction

16. Endothermic Reaction increases

17. Some rules for heat and work Heat given off is negative. (Exothermic) Heat absorbed is positive. (Endothermic) Work done by system on surroundings, heat flows out of system, heat is product. Work done on system by surroundings, heat flows into system, heat is reactant. Thermodynamics- The study of energy and the changes it undergoes.

18. System View Heat q Work w Endothermic (Heat Reactant) ++ Exothermic (Heat Product) -- Take The Systems Point Of View

19. Calculate ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system. Time to plug and chug Remember to watch signs

20. Piston System & Surroundings

21. Work Done By A Piston 1.Since pressure is defined as force per unit area pressure of gas is 2.Work is defined as force applied over a distance and piston moves ∆ h 3.Work = force x distance = F x ∆ h 4.Use equations in 1 & 3, then w = P x A x ∆ h 5.Since Volume = A x ∆ h = ∆V can be substituted into work in step 4 and get

22. Work Done By A Piston If gas is expanding it is work done on the surroundings and work has to be a negative number and ΔV is positive (V final – V initial ) w and PΔV must have opposite signs Hence the negative in formula Calculate the work associated with expansion of a gas from 46 L to 64 L at a constant external pressure of 15 atm.

23. Practice A balloon is being inflated to its full extent by heating the air inside it. In the final states of this process, the volume of the balloon changes from 4.00 x 10 6 L to 4.50 x 10 6 L by the addition of 1.3 x 10 8 J of energy as heat. Assuming that the balloon expands against a constant pressure of 1.0 atm, calculate ∆E for the process. Work involved in the expansion or compression of gases is called pressure-volume work. Equality you should know:

24. Enthalpy (H) A measure of total energy in system Enthalpy is defined as H = E + PV When a change occurs at constant pressure the equation becomes: Δ H = Δ  E + PV) = Δ E + P∆V the heat at constant pressure q p can be calculated Δ E = q p + w and w= – P Δ V Solve for q p q p = Δ E + P Δ V = Δ H End result of calculation: q p =  Δ H

25. Enthalpy (H) For a chemical reaction, enthalpy change is ∆H = H products – H reactants If H products > H reactants +∆H Endothermic If H reactants > H products -∆H Exothermic Equality you should know: 1 mole X = (value) ∆H x kJ

26. Determine the Sign of ∆H Indicate the sign of ∆H in each of the following process carried out under atmospheric pressure and indicate if the process is endothermic or exothermic: 1.An ice cube melts 2.1 g of butane (What is the formula?) is combusted in sufficient O 2 to give complete

27. Calorimetry Device used to measure heat associated with a chemical reaction is a calorimeter Calorimetry is the Science Of Measuring Heat Two kinds of devices: – Constant Pressure Calorimeter – Coffee cup calorimeter – Bomb Calorimeter – constant volume heat capacity ( C ) for a substance The heat capacity of an object is the amount of heat required to raise its temperature by 1 K

28. Calorimetry Specific Heat Capacity heat capacity is given as per gram of substance Units are J/ o C * g or J/K * g Molar Heat Capacity heat capacity is given as per mole of substance Units are J/ o C * mol or J/K * mol Basic Principle behind the Calorimeter is Energy Released by Reaction (Joules) q = Energy Absorbed by Solution q = mass of solution (m) * Specific heat capacity (C) * Change in temperature ( ∆T) q = m * C * ∆T

29. Practice Constant Pressure Calorimeter 1.How much heat is needed to warm 250g of water from 22 o C to 98 o C? The specific heat of water is 4.18J/gK. 2.What is the molar heat capacity of water?

30. Practice Constant Pressure Calorimeter A hot water bottle filled with 0.75kg of water is at 80.0 O C at the start of the night and cools to 20.0 O C by morning. How much heat was given out? CH 2 O (l) = 4180 J/kg O C

31. 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 calorimeters are known and tested Δ E = q + w, and and V=0, then w=0 and Δ E = q in a constant volume process. q rxn = -C calorimeter x ∆T

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

33. Measuring q rxn Using Bomb Calorimeter Methylhydrazine (CH 6 N 2 ) is commonly used as a liquid rocket fuel. Yea – finally some rocket science! The combustion of methylhydrazine with O 2 produces N 2(g), CO 2(g), and H 2 O (l) : CH 6 N 2 + O 2 → N 2(g) + CO 2(g) + H 2 O (l) When 4.00 g of methylhydrazine is combusted in a bomb calorimeter, the temperature of the calorimeter increases from 25.00 O C to 39.50 O C. In a separate experiment the heat capacity of the calorimeter is measured to be 7.794kJ/ O C. What is the heat of reaction for the combustion of a mole of CH 6 N 2 in this calorimeter?

34. Properties Intensive Properties not related to the amount of substance. – density, melting point, temperature. Extensive Properties - does depend on the amount of stuff. – Heat capacity (C), mass, heat of reaction (q), volume

35. Hess’s Law Enthalpy is a state function. It is independent of the path. We can add equations to come up with the desired final product, and find Δ H Two rules If the reaction is reversed the sign of Δ H is changed If the reaction is multiplied, so is Δ H Calculate the Δ H 1 for the overall reaction

36. N2N2 2O 2 O2O2 NO 68 kJ NO 2 ∆H 2 = 180 kJ ∆H 3 = -112 kJ H (kJ) ∆H 1 = ∆H 2 + ∆H 3 Hess’s Law Hess’s Law

37. Net Reaction achieved by summing the 2 reactions Enthalpy achieved by summing enthalpy for 2 step reactions

38. Practice Two forms of carbon are graphite, the soft, black, slippery material used in “lead” pencils and as a lubricant for locks, and diamonds, the brilliant, hard gemstone that is a lady’s best friend. Using the enthalpies of combustion for graphite (-394kJ/mol) and diamond (-396kJ/mol), calculate ∆H for the conversion of graphite to diamond: C graphite (s)  C diamond (s) Two combustion reactions for the forms:

39. Reverse the 2 nd equation and change sign of ∆H Do #58 & 60 on p. 269

40. Standard Enthalpy Standard Enthalpy of Formation ( Δ H f o ) is the change in enthalpy that accompanies the formation of one mole of a pure substance from its elements with all substances in their standard state. Degree symbol on Δ H f o indicates standard states:  Δ H f o For Compounds:  Gaseous substance, pressure is exactly 1 atm  For liquid or solid it is the pure liquid or pure solid  For substance in solution, it is the concentration of 1M  Δ H f o for an element: it is the form the element exists under 1 atm and 25 o C, Ex. oxygen is O 2 (g) at 1 atm

41. Standard Enthalpies of Formation Enthalpy is a state function so could use Hess’s Law Choose any pathway from reactant to product and sum the enthalpy changes along the chosen pathway. Standard states are 1 atm, 1M and 25 o C For an element Δ H f o = 0, there is no formation reaction needed when element is already in its standard state See table in Appendix 4 (p. A19-22)

42. How can you find Δ H w/out Hess’s Law? We can use Heats of Formation Δ H f o to figure out the heat of reaction. Ammonia is burned in air to form nitrogen dioxide and water – Always start with a balance equation 4NH 3(g) + 7O 2(g)  4NO 2(g) + 6H 2 O (l) And use this formula Do #68 p. 269

43. Other Enthalpies Enthalpies of vaporization - ∆H for converting a liquid to a gas Enthalpies of fusion - ∆H for melting solids Enthalpies of combustion - ∆H for combusting a substance in oxygen