1. Chapter 16 – Reaction Energy
Thermochemistry – Study of the transfers of energy as heat, that accompany chemical and physical changes.

2. Heat and Temperature
All matter is in constant, random motion, regardless of temperature. (KMT)
Only temperature in which there is no kinetic energy is at absolute zero.
-273 oC = 0 kelvins
In theory, the temperature at which all kinetic energy equals zero.
Heat is a result of this constant motion.
More heat = More motion

3. Heat and Temperature Heat (Heat Energy) Temperature
Sum total of the kinetic energy (KE) of the particles in a sample of matter.
Flows from high temp objects to low temp objects.
Calorimeter – device used to measure heat loss or gained by a sample of matter.
Temperature
Measures the average KE of particles in a sample of matter. (The degree of how hot or cold something is.)

4. Standard
Calorimeter

5. Bomb
Calorimeter

6. Units of Heat (q) Joule (J) – S.I. Unit of heat energy.
calorie (cal) – quantity of heat required to raise the temperature of 1 g of water by 1oC.
1 cal = 4.184J
Calorie (Cal) – Dietary calorie.
1Cal = 1000cal

7. Converting units of energy
1 cal = J
1 Cal = 1000 cal
1 kJ = 1000 J
How many kJ are in 2.45 x 108 cal?

8. Practice: 1) Convert the following to calories (cal): a) 900 J
2) Convert the following to joules (J):
a) 220 Cal

9. Heat Capacity and Specific Heat
Heat Capacity – amount of heat needed to raise the temperature of a sample 1oC.
Units (J/oC)
Will change depending upon the mass of the object.
More mass more heat capacity.

10. Specific Heat Capacity
Amount of heat needed to raise the temperature of 1g of a substance by 1oC.
Larger value = Longer time to Heat
Specific Heat of Water
4.184 J/goC = 1 cal/goC

11. Substance
Specific Heat(j/goC)
Water (liquid)
4.184
Carbon
(graphite)
.720
Water (ice)
2.092
Carbon (diamond)
.502
Water (steam)
2.013
Iron
.444
Ethyl Alcohol
2.452
Acetic Acid
2.048
Sugar
.899
Copper
.385
Silver
.237
Lead
.129

12. Heat Capacity vs. Specific Heat Capacity
Aluminum has a specific heat capacity of.899 J/goC.
How much heat will it take to raise 1 g of Al by 1oC?
Will 500g of Al or 100g have a greater heat capacity?
Will 500g of Al or 100g have a greater specific heat capacity?

13. Specific Heat Equation
Using specific heat capacity
Ti = Initial Temperature
Tf = Final Temperature
∆T = Change in Temperature
q = heat
m = mass
Cp = Specific Heat Capacity

14. Sample
What is the specific heat of a 550g substance, that gains 1200J of heat, as it increases in temperature from 25oC to 40oC?

15. Sample 2
How much heat is needed to change the temperature of 150g of Aluminum from 20oC to 100oC?
Know the Specific Heat of Al = .899J/goC

16. Calorimetry Problem
In using a calorimeter that contains 500.g of water, 100.g of an unknown metal at 120.oC is added to the water. The water increases in temperature from 25.0oC to 28.0oC. What is the specific heat of the unknown metal?

17. Heating Curves As an object gains heat:
It may increase the objects temperature
It may cause a phase change
Or may cause both a phase change and a change in temperature.
Temperature changes can be measured by:
q = Cp x m x T
Phase changes can be measured by:
Molar heat conversions:
40.7 kJ/mol or 6.01 kJ/mol

18. Sketch the heating curve for 50g water as it starts at –20oC and increases to 60oC.
Liquid
melting
0oC
Ice
-20oC
Heat Increasing 

19. Sample Problem
Determine the heat gained by the 50g of water as it increases in temperature from –20oC to 60oC.
60oC
Liquid
melting
0oC
Ice
-20oC
Heat Increasing 
Heat at the point of Ice ( q = Cp x m x T) Cp = 2.092J/goC
Heat at the melting phase (Molar Heat of Fusion)
Heat at the point of Liquid (q = Cp x m x T)

20. Calculations

21. Practice
Determine the energy needed to change 150g of water from 35oC to 110oC. 1) 2) 3)

22. Heat Summary Summary of Heat Calorimetry
Heat gained by one object = Heat lost by another.
Heat gained = Heat Lost
If one object gains 1250J then another object must lose 1250J. (1250J = -1250J)
Calorimetry
Study of heat flow and heat measurement.
Usually involves the temperature of water.

23. Heat of Solutions
When a solute dissolves in a solution the solution always has a enthalpy change.
NaOH in a solution is very exothermic.
NH4NO3 in a solution is endothermic.
Similar solution in cold packs.
Determine heat of solution by using calorimeter.

24. Practice Heat of Solution
When 20.0g sample of KOH dissolves in 50g of water in a calorimeter, the temperature increases from 18oC to 32oC. Calculate the H (kJ/mol) for the process.
KOH  K OH –1
q = Cp  m  T
q =

25. Problem:
When 12.8g sample of KCl dissolves in 75g of water in a calorimeter, the temperature decreases from 32oC to 22.6oC. Calculate the H (kJ/mol) for the process.
KCl  K Cl –1
q =

26. Summary of Heat and Physical Change
When substances are changing physically (phase change or solutions):
1) You can use q = Cp x m x T
When there is a temperature change.
2) You can convert a given mass into moles and use a molar heat of fusion or vaporization.
When there is a phase change.
Melting = 6.01 kJ/mol
Boiling = 40.7 kJ/mol

27. Heat Transfer System Surroundings Universe
Substance being observed or object in which all focus is on.
Surroundings
Everything outside the object being studied.
Universe
System and Surroundings combined.

28. Types of Heat Reactions
Exothermic
Heat is transferred from the system to the surroundings.
Object (system) will decrease in temperature.
Endothermic
Heat is transferred from the surroundings to the system.
Object (system) will increase in temperature.

29. What is exothermic and what is endothermic?

30. Thermochemical Reactions
Exothermic reactions have heat in the product side of the equation.
Ex. Water  Ice + Heat
C3H8 + 5O2  3CO2 + 4H2O kJ
All Combustion reactions.
Endothermic reactions have heat in the reactant side of the equation.
Ex. Ice + Heat  Water
C + H2O + 113kJ  CO + H2

31. Heat of Reaction(Enthalpy)
Enthalpy (H)
The heat of a reaction at a constant pressure.
Enthalpy Change (H)
The change in enthalpy of a reaction is equal to the heat gained or absorbed during the reaction.
H reaction = Hproducts - Hreactants
H reaction = (+), endothermic
H reaction = (), exothermic

32. Standard Enthalpy (Ho)
Enthalpy change that is measured when the reactants in their standard states, change into their products at standard states.
Express the following equation using energy if H = kJ
C3H8 + 5O2  3CO2 + 4H2O

33. Use of Enthalpy
The heat of a reaction, is the energy in kilojoules, equivalent to the mole ratios of reactants and products.
C3H8 + 5O2  3CO2 + 4H2O kJ
1 mol C3H8 = 2043kJ of heat
5 mol O2 = 2043kJ of heat
3 mol CO2 = 2043kJ of heat
4 mol H2O = 2043kJ of heat

34. Heat of Reaction Problems
Determine the mass of oxygen needed to produce 50,000kJ of energy.
C3H8 + 5O2  3CO2 + 4H2O kJ
How many kilojoules of heat would be produced if 55g of C3H8 reacted?

35. Hess’s Law
When a series of reactions are added together, the enthalpy change for the overall reaction will be the sum of the individual enthalpy changes of each reaction.

36. Applying Hess’s Law
Match the products and reactants in the individual reactions to the net reaction.
Step 1
If a compound needs to be on the other side of the equation, then flip the entire reaction. When flipping a reaction, change the sign on the H.
Step 2
If the reaction has to be multiplied by a coefficient then, multiply the H by that same multiple.

37. Applying Hess’s Law
Determine the enthalpy change for the following equation:
N2 + 2O2  2NO2 H = ?
Use the following equations:

38. Solving the Problem

39. Practice Problem: Hess’s Law
Using the following reactions and their enthalpy changes:
2SO2 + O2  2SO3 H0 = -196kJ
2S + 3O2  2SO3 H0 = -790kJ
Calculate the standard enthalpy change for the combustion of sulfur to produce SO2.
S + O2  SO2 H0 = ?

40. Practice Answer

41. Heats of Summation
Alternative method of determining the change in enthalpy of a reaction.
H0 = (Products) - (Reactants)
Remember to multiply the number of moles of each substance by the heat of formation.

42. Example Heat of Summation
Determine the change in enthalpy for the following equation:
2SO3(g)  2SO2(g) + O2(g)
Use table A.6 for the standard heats of formation values.

43. Practice Heat of Summation
Determine the enthalpy change for the following equation:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)

44. 16.2 Driving Force of Reactions
Entropy (S)
Amount of disorder in a system.
Measure the degree of randomness of the particles in a system.
units: kJ/(mol●K)
“Law of Entropy” or the “Law of Disorder”
States that things move in the direction of maximum disorder or chaos.
Think of your room!

45. States of Matter
Solids < Liquids < Gases

46. Number of Molecules
Entropy increases whenever reaction gives increase in number of gaseous particles.

47. Molecular Size
For same type of bonding, larger molecules possess higher entropy.
Relative Entropies:
CH4 < C2H6 < C3H8 < C4H10

48. Mixing

49. Expansion

50. Temperature
Entropy increases with increasing temperature (kinetic energy).
Equilibrium shifts towards the higher entropy!

51. Entropy of a Reaction
Which side of the reaction has the greatest entropy?
4NO(g) +6H2O(g) → 4NH3(g) + 5O2(g)
2NaCHCO3(s)→Na2CO3(s) +CO2(g) + H2O(g)

52. Free Energy (G) Energy that is available to do work.
The combined function of the enthalpy-entropy process determines whether a reaction will proceed spontaneously at room standard conditions.
∆Go = ∆Ho - T∆So
∆Go= (-), Spontaneous
∆Go= (+), Non-spontaneous

53. Spontaneous Reactions
Determining whether a reaction will occur without any additional energy at standard condition is termed spontaneous.
∆Ho
∆So
Spontaneous
(-)
(+)
(Low Temperature)
(High Temperature)
Non-spontaneous

54. Sample
CaCO3  CaO + CO2
Given that this reaction takes place at 20oC. What is the free energy of this reaction? Is it spontaneous or not?
∆Ho = kJ/mol ∆So = J/K●mol.