1. Thermochemistry
Internal Energy
Kinetic energy
Potential energy

2. Thermochemistry
Internal Energy
Kinetic energy
Potential energy

4. Chemical Energy Changes
System and Surroundings
Exothermic Reaction

5. Endothermic Reaction

6. Thermochemistry
Thermochemisty is the study of the relationship between heat and chemical reactions.
1. Kinetic energy is energy possessed by matter because it is in motion
Thermal energy-- random motion of the particles in any sample above 0 K
Heat -- causes a change in the thermal energy of a sample. Flows from hot to cold

7. Heat

8. Potential Energy
2. Potential energy is energy possessed by matter because of its position or condition.
A brick on top of a building has potential energy that is converted to kinetic energy when it is dropped on your head
Chemical energy is energy possessed by atoms as a result of forces which hold the atoms together (Boxes!)

9. Where is the Energy? Definitions we will use: System: Reaction (bonds)
Surrounding: solvent, reaction
vessel, air, etc.
An everyday example: burning wood
Initially, much energy stored as potential in C-H bonds, little kinetic energy in the air
Finally, lower potential energy in the C=O bonds, higher kinetic energy in the air

10. Total Energy Total Energy = kinetic + potential
Law of Conservation of Energy - The total energy of universe is constant
Internal Energy - E - the sum of all the kinetic and potential energies of all the atoms and molecules in a sample.

11. Change in Energy of System
Change in internal energy of system = heat + work
Convention: point of view of system

12. Change in Internal Energy
DE = q + w
Work = Force x distance
What happens to your internal energy when you push a boulder?
What happens to your internal energy when you push a boulder on a rough surface?

13. Chemical Work ׀W׀ = ׀F x Dh׀ P = F/A ׀W׀ = ׀P x A x h׀ ׀W׀ = ׀PV׀
Sign Convention:
W = -PDV

14. Test Your Understanding
For the following three reactions:
Are they performed under constant pressure or not?
What is the sign of work in each case?

15. State Function State Function Path Dependent Internal Energy Pressure
Volume
Path Dependent
Work
heat
Property depends only on present state

16. Enthalpy Most reactions are done in open containers, so P is constant
Need a term for constant pressure where only work is PV
At constant pressure, qp
ΔE = qp – PΔV
qp = ΔE + PV
ΔH = ΔE + PΔV (definition)
ΔH = qp

17. Enthalpy IF pressure is constant and only work is PV
Change in enthalpy is equal to flow of energy in form of heat
Measurable by Temperature
Change in enthalpy is “heat of reaction”
If no net change in moles of gas, enthalpy ~ energy

18. Enthalpy Signs on ΔH Endothermic ΔH = + Exothermic ΔH = -
+ heat is taken in by system
- heat is given off by system
Endothermic ΔH = +
Exothermic ΔH = -

19. Exothermic
2 Al (s) + Fe2O3 (s)  Al2O3 (s) Fe (s) + energy
Which bonds have more potential energy? Which bonds are stronger?

20. Endothermic Ba(OH)2. 8H2O (s) + 2 NH4SCN (s) + energy
 Ba(SCN)2 (aq) NH3 (g) H2O (l)

21. Reaction Enthalpy It is useful to know how much energy is released
CH4 (g) O2 (g)  CO2 (g) H2O (l) + ENERGY
It is useful to know how much energy is released
CH4 (g) O2 (g)  CO2 (g) H2O (l) ΔH = kJ
This is a stoichiometric amount
890 kJ released = 1 mol CH4 = 2 mol O2 = 1 mol CO2 = 2 mol H2O

22. Reaction Enthalpy Depends on coefficients, direction, and phases
2 CH4 (g) + 4 O2 (g)  2 CO2 (g) + 4 H2O (l) ΔH = kJ
CO2 (g) H2O (l)  CH4 (g) O2 (g) ΔH = kJ
CH4 (g) O2 (g)  CO2 (g) H2O (g) ΔH = -802 kJ

23. Reaction Enthalpy

24. Standard Reaction Enthalpy
Standard State - a compound in its pure state at 1 atm pressure, all solutions are 1 M
Temperature can vary but usually K
Standard Reaction Enthalpy (ΔHro)- reaction enthalpy when all products and reactants are in the standard state
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l) ΔHo = kJ

25. Enthalpy and Stiochiometry

26. Application: Enthalpies of Combustion
Standard Enthalpy of Combustion (ΔHco) -- change in enthalpy for the combustion of one mole substance at standard conditions
Combustion is combination with O2 to give CO2 and water
Specific Enthalpy -- the enthalpy of combustion per gram
enthalpy density -- enthalpy of combustion per liter

27. Enthalpies of Combustion

30. How Do We Determine Standard Reaction Enthalpies?
Tabular Data: Hess’s Law
Combining appropriate reactions
Heat of Formation
Bond enthalpies
Experimental
Calorimetry: constant pressure
Calorimetry: constant volume

31. Hess’s Law
Just a restatement of the first law

32. How Do We Determine Standard Reaction Enthalpies?
Tabular Data: Hess’s Law
Combining appropriate reactions
Heat of Formation
Bond enthalpies
Experimental
Calorimetry: constant pressure
Calorimetry: constant volume

34. Using Hess’s Law Find ΔHo for C (s) + ½ O2 (g)  CO (g)
C (s) + O2 (g)  CO2 (g) ΔHo = kJ
2 CO (g) + O2 (g)  2 CO2 (g) ΔHo = kJ

35. How Do We Determine Standard Reaction Enthalpies?
Tabular Data: Hess’s Law
Combining appropriate reactions
Heat of Formation
Bond enthalpies
Experimental
Calorimetry: constant pressure
Calorimetry: constant volume

36. Standard Enthalpies of Formation
The standard enthalpy of formation is the enthalpy change when one mole of a substance in its standard state is formed from the elements in their standard states. Hof
Write an equation for the standard heat of formation of carbon dioxide
C(s) O2 (g)  CO2 (g) Hof (CO2)
Write an equation for DHof of CH3OH
Write an equation for DHof of N2 (g) and explain why its value is zero.

37. Standard Enthalpies of Formation
DHof can be compiled in table form

38. Application of Heat of Formation
Hess’s Law
Without the limitations of combining limited number of reactions
Common starting point: elements

39. Calculating Heat of Reaction
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l)
ΔHor = ΣHof (products) - ΣHof (reactants)

40. How Do We Determine Standard Reaction Enthalpies?
Tabular Data: Hess’s Law
Combining appropriate reactions
Heat of Formation
Bond enthalpies
Experimental
Calorimetry: constant pressure
Calorimetry: constant volume

41. Using Bond Enthalpies Most versatile Least exact
Must be able to draw Lewis Dot structures

42. Bond Enthalpies Bond Dissociation Energies Hess’s Law
Positive values
Hess’s Law
In principle, “free atoms” formed
DHrxn = SBDE(broken) – SBDE(formed)

43. Bond Enthalpies
Calculate the heat of reaction for the combustion of formamide (CH3NO).
Equation
Lewis Dot
Calculate

44. How Do We Determine Standard Reaction Enthalpies?
Tabular Data: Hess’s Law
Combining appropriate reactions
Heat of Formation
Bond enthalpies
Experimental
Calorimetry: constant pressure
Calorimetry: constant volume

45. Heat Capacity
Heat Capacity (C) - the heat required to raise the temperature of an object by 1 K
C = q/ΔT
extensive property

46. Specific Heat Capacity
Specific heat capacity (Cs) - the heat required to raise 1 g of a substance by 1 K
Specific heat
Cs = C/m
intensive property
q = m Cs ΔT

48. Each Substance “Stores” Heat Differently
Increased
263 oC
25 oC
1 g copper
100 J
Increased
24 oC
25 oC
1 g water
Putting the same amount of heat into two different substances will raise their temperature differently based on the specific heat of each substance
q=mCsDT

49. Calorimetry
Calorimeter - an insulated container fitted with a thermometer
Open to atmosphere, so P is constant
qp = m Cs ΔT
qp = ΔH

50. Calorimetry Problem
In a coffee cup calorimeter, 50.0 mL of M silver nitrate and 50.0 mL of M HCl are mixed. The following reaction occurs: Ag+ (aq) + Cl- (aq)  AgCl (s). If the two solutions are initially at oC, and if the final temperature is oC, calculate the change in enthalpy for the reaction. (What assumptions need to be made?)

51. Bomb Calorimeter Volume is constant q = ΔE qsystem= -qsurroundings
qrxn= -(qbomb+ qwater)
qwater = mwaterCsΔT
qbomb = mbombCsΔT or CDT

52. Thermodynamics of Ideal Gas
Heat capacity of monoatomic gas
From KMT, KE = 3/2RT
KE = translational energy
Heat required to raise temp 1 degree is 3/2R
At constant volume, no work is done
Molar heat capacity Cv = 3/2R = J/mol K
Is this constant for all gases?

53. Consider polyatomic gases
For polyatomic molecules, _____ energy has to be put into the same amount of gas to raise it by 1 degree
Where does this energy go? (Not translational!)
Trend??

54. Monoatomic Gas at Constant Pressure
Energy input does two things: increase translational energy (T) and expand gas (w)
w = PDV = nRDT
Molar heat capacity
Cp = 3/2R + R = 5/2R

55. Polyatomic at Constant Pressure
Explain physical basis of this equation:
Cp =Cv + R
This is observed!
When would you see a deviation from Cp-Cv = R?

56. Summary

57. Conceptual Understanding of Gas Cycle

58. Thermochemistry of Physical Change
Vaporization - endothermic process
Vapor has higher H that a liquid at the same temperature
Enthalpy of vaporization ΔHvap --
ΔHvap = Hvapor - Hliquid

59. Freezing, Melting and Sublimation
Enthalpy of fusion ΔHfus (melting)
ΔHfus = Hliquid - Hsolid
Enthalpy of freezing = - ΔHfus
Enthalpy of sublimation, ΔHsub
ΔHsub = Hvapor - Hsolid

62. Heating/Cooling Curve

63. Enthalpy of Solution
Enthalpy (Heat) of Solution – ΔHsoln – heat change associated with the dissolution of a known amount of solute in a known amount of solvent.
ΔHsoln = Hsoln – Hcomponents

65. Lattice Energy
STEP 1
Lattice Energy (U) – the energy required to separate 1 mole of a solid ionic compound into gaseous ions.
NaCl (s)  Na+(g) + Cl-(g)
U = 788 kJ/mol

66. Heat of Hydration
Heat of hydration – ΔHhydr – Enthalpy change associated with the hydration process.
Na+(g) + Cl-(g)  Na+(aq) + Cl-(aq)
ΔHhydr = -784 kJ

67. ΔHsoln= U +ΔHhydr= 788 kJ/mol + (-784 kJ) = 4 kJ/mol