1. Energy and Chemical Reactions
Chapter 5

2. Learning Objectives Students understand
And apply the first law of thermodynamics

3. Learning Objectives Students will be able to
Assess the transfer of energy as heat
Calculate P-V work done by a system and use energy as heat and work to calculate changes in internal energy
Calculate energy and enthalpy changes from calorimetry data
Calculate energy evolved or required using thermodynamic data

4. 5.1 Energy the science of heat and work is called thermodynamics
The kinetic energy of molecules is referred to as thermal energy.

5. Energy
The Law of Conservation of Energy states that energy can be converted from one form to another, but can neither be created nor destroyed
restated mathematically as the first law of thermodynamics

6. Systems and Surroundings
A system is an object, or collection of objects, being studied.
The surroundings include everything outside the system that can exchange energy and/or matter with the system.

7. Temperature and Heat heat is not the same as temperature
more thermal energy = more particle motion
total thermal energy is sum of individual energies of all particles

8. Thermal Equilibrium
thermal equilibrium means that two objects have reached the same temperature
heat always transfers from hotter object to cooler object
transfer continues until equilibrium is reached
heat lost = heat gained

9. Thermal Equilibrium
exothermic process occurs when heat transfers from system to surroundings
heat is released as a product
qsys < 0
endothermic process occurs when heat transfers from surroundings into system
heat is required as a reactant
qsys > 0

10. 5.2 Specific Heat quantity of energy transferred depends on
Quantity of material
Magnitude of the temperature change
Identity, including phase, of the material gaining or losing energy

11. 5.2 Specific Heat
specific heat is the quantity of heat required to raise 1 gram of substance 1 kelvin (J/g∙K)
q = CmΔT
Where ΔT = Tfinal - Tinitial

12. Practice Problem
In an experiment it was determined that 59.8 J was required to change the temperature of 25.0 g of ethylene glycol by 1.00 K. Calculate the specific heat capacity of ethylene glycol.

13. Quantitative Aspects
In an isolated system, the sum of the energy changes must be zero!
Energy is CONSERVED!!

14. Practice Problem
A 15.5 g piece of chromium, heated to oC, is dropped into 55.5 g of water at 16.5 oC. The final temperature of the metal and the water is 18.9 oC. What is the specific heat of chromium? Assume no heat is lost to container or air.

15. Homework
After reading sections 5.1 and 5.2, you should be able to do the following problems…
p. 217 (3-4, 9-16)

16. 5.3 Energy and Changes of State
Heat of fusion: solid to liquid (endothermic)
Heat of vaporization: liquid to gas (endothermic)
Heat of solidification: liquid to solid (exothermic)
Heat of condensation: gas to liquid (exothermic)

17. Energy and Changes of State
Temperature is constant during a state change!!

18. Practice Problem
How much heat must be absorbed to warm 25.0 g of liquid methanol, CH3OH, from 25.0 oC to 64.6 oC and then to evaporate the methanol completely at that temperature? The specific heat of methanol is 2.53 J/g∙K. ΔHvap = 2.00 x 103 J/g

19. Practice Problem
To make a glass of iced tea, you pour 250 mL of tea, whose temperature is 18.2 oC, into a glass containing 5 ice cubes. Each cube has a mass of about 15 g. What quantity of ice will melt, and how much ice will remain to float at the surface in this beverage? Iced tea has a density of 1.0 g/cm3 and a specific heat of 4.2 J/g∙K. Assume no energy transferred between system and surroundings.

20. 5.4 First Law of Thermodynamics
the energy change for a system is the sum of the energy transferred as heat and the energy transferred as work
ΔU = q + w
Where
ΔU is change in internal energy
q is heat transfer to or from system
w is work transfer to or from system, called P-V (pressure-volume) work

21. P-V work
work done on or by the system equals the volume change that occurs against resisting external pressure
w = - P(ΔV)
If energy is transferred at a constant volume process, energy transferred as work is zero and the ΔU = q

22. Enthalpy Enthalpy is heat content of a substance at constant pressure
Enthalpy change (ΔH)
negative ΔH and ΔU mean that energy is transferred from system to surroundings
positive signs mean that energy is transferred from surroundings to system

23. State Functions
changes in ΔH or ΔU depend only on final and initial states
the pathway to go between is not relevent (example: bank accounts)

24. Homework
After reading sections 5.3 and 5.4, you should be able to do the following problems…
p. 217a (21-30)

25. 5.5 Enthalpy Changes
enthalpy changes are specific to the reactants and products and their amounts (states of matter are important)
ΔH depends on the number of moles of reaction as written
ΔH is positive when endothermic and negative when exothermic (reverse reactions have opposite sign)

26. Enthalpy Changes
Standard reaction enthalpy: enthalpy change accompanying a specific reaction
Standard state: most stable form of the substance in the physical state at a given pressure and temperature

27. Enthalpy Changes change depends on molar amounts of substances
calculate moles of substance and then multiply that by heat transfer per mole

28. Practice Problem
What quantity of heat energy is required to decompose 12.6 g of liquid water to the elements?
The combustion of ethane, C2H6, has an enthalpy change of kJ for the reaction. Calculate ΔH when 15.0g is burned.
2C2H6(g) + 7O2(g)  4CO2(g) + 6H2O(g)

29. 5.6 Calorimetry technique to determine heat transfer
constant pressure measures change in enthalpy
constant volume measures change in internal energy

30. Constant Pressure Calorimetry
can be used to determine heat gained or lost by solution
can be used to determine heat required or released by reaction
at constant pressure, the heat measured is ΔH
change in heat content of solution can be measure and used to calculate the heat of reaction
qrxn + qsoln = 0

31. Constant-Pressure Calorimetry

32. Practice Problem
Assume you mix 200. mL of M HCl with 200. mL of M NaOH in a coffee-cup calorimeter. The temperature of the solutions before mixing was 25.20oC; after mixing and allowing reaction to occur, the temperature is 27.78oC. What is the molar enthalpy of neutralization of the acid? (assume densities of all solutions are 1.00 g/mL and their specific heats are 4.20 J/gK)

33. Constant Volume Calorimetry
used to calculate heats of combustion and caloric value of foods – use a “bomb”
constant volume, so energy transfer as work doesn’t occur and heat is therefore the change in internal energy
qr + qbomb + qwater = 0

34. Constant Volume Calorimetry

35. Practice Problem
A 1.00g sample of sucrose is burned in a bomb calorimeter. The temperature of 1.50 x 103 g of water in the calorimeter rises from oC to oC. The heat capacity of the bomb is 837 J/K and the specific heat of water is 4.20 J/gK. Calculate (a) the heat evolved per gram of sucrose and (b) the heat evolved per mole of sucrose.

36. Homework
After reading sections 5.5 and 5.6, you should be able to do the following…
p. 217a-b (33-42)

37. 5.7 Hess’s Law
sometimes products immediately undergo other reactions and therefore calorimetry cannot be used
Hess’s Law states that if a reaction is the sum of two or more other reactions, ΔH for the overall process is the sum of the ΔH values of those reactions

38. Energy Level Diagrams
can visualize Hess’s Law by using diagrams to show the formation of different steps in a reaction as well as the enthalpy changes involved

39. Energy Level Diagrams

40. Practice Problem
What is the enthalpy change for the formation of ethane, C2H6, from elemental carbon and hydrogen?
2C(s) + 3H2(g)  C2H6(g)

41. Practice Problem
Use Hess’s Law to calculate enthalpy change for the formation of CS2(l) from C(s) and S(s) from the following enthalpy values.
C(s) + O2(g)  CO2(g) ΔH = kJ
S(s) + O2(g)  SO2(g) ΔH = kJ
CS2(g) + 3O2(g)  CO2(g) + 2SO2(g) ΔH = kJ
C(s) + 2S(s)  CS2(g) ΔH = ?

42. Standard Enthalpies
Standard molar enthalpy of formation, ΔHof, is the heat change for the formation of 1 mol of a compound directly from its component elements in their standard states.
See appendix L

43. Standard Enthalpies
Standard enthalpy of formation for an element in its standard state is zero.
Most values are negative
can compare thermal stability; more exothermic is more stable

44. Enthalpy Change for a Reaction
ΔHorxn = Σ[ΔHof(products)] - Σ[ΔHof(reactants)]
Find enthalpies in a table and then plug the values into the equation above

45. Practice Problem
Calculate the standard enthalpy of combustion for benzene, C6H6.
C6H6(l) + 7½O2(g)  6CO2(g) + 3H2O(l)
ΔHof[C6H6(l)] = kJ/mol

46. Practice Problem
Case Study p. 212: The Fuel Controversy – Alcohol and Gasoline

47. 5.8 Favored Reactions product-favored reactions go from left  right
most reactions that are exothermic (have negative values of enthalpy) are product-favored
most reactions that are endothermic (have positive enthalpy change) are reactant-favored

48. Practice Problem
Calculate ΔHorxn for each of the following reactions and decide if the reaction may be product- or reactant-favored.
2HBr(g)  H2(g) + Br2(g)
C(diamond)  C(graphite)

49. Homework
After reading sections 5.7 and 5.8, you should be able to do the following problems…
p. 217c-d (51-52,63-64,115)