1. Chapter 5 “Thermochemistry”
Mr Daniel : OHS AP Chemistry
Credit to:
Charles Page High School
Stephen L. Cotton

2. Section 11.1 The Flow of Energy – Heat and Work
OBJECTIVES:
Explain how energy, heat, and work are related.
Classify processes as either exothermic or endothermic.
Identify the units used to measure heat transfer
Distinguish between heat capacity and specific heat.

3. Energy Transformations
Thermochemistry – study of heat changes that occur during chemical and physical changes.
Energy - capacity for doing work or supplying heat
weightless, odorless, tasteless
Energy stored within chemical substances is called chemical potential energy
Gasoline contains a significant amount of chemical potential energy

4. Energy Transformations
Heat (“q”) is energy that transfers from one object to another, because of a temperature difference between them.
only changes of heat can be detected!
flows from warmer  cooler object
All chemical reactions and changes in physical
state involve either:
release of heat, or
absorption of heat

5. Exothermic and Endothermic Processes
When studying heat changes there are two important categories:
the system - the part of the universe on which you focus your attention (a test tube, a room, a molecule, the Earth etc.)
the surroundings - includes everything else in the universe
Together, the system and it’s surroundings constitute the universe

6. Exothermic and Endothermic Processes
Thermochemistry is concerned with the flow of heat from the system to it’s surroundings, and vice-versa.
The Law of Conservation of Energy states that in any chemical or physical process, energy is neither created nor destroyed.
All the energy is accounted
for as work, stored energy,
or heat.

7. Exothermic and Endothermic Processes
Endothermic: Heat flowing into a system from it’s surroundings
defined as positive change.
q has a positive value
The system gains heat as the surroundings cool down

8. Exothermic and Endothermic Processes
Exothermic: Heat flowing out of a system into
it’s surroundings
defined as negative change
q has a negative value
The system loses heat
as the surroundings heat up

9. Exothermic and Endothermic
Every reaction has an energy change associated with it. For example, the…
** Gummy Bear Sacrifice!! ** (next week!!!)
Exothermic reactions release energy, usually in the form of heat.
Endothermic reactions absorb energy
Chemical Energy is stored
in the form of bonds
between atoms.

10. Units for Measuring Heat Flow
A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1 oC.
a Calorie, (written with a capital C), is calories and refers to the energy in food
(100 Calorie candy bar is 100,000 calories)
1 Calorie = 1 kilocalorie = 1000 calories

11. Units for Measuring Heat Flow
The Joule is the SI unit of heat and energy and is related to the calorie:
4.184 J = 1 cal
Named after James Prescott Joule

12. Heat Capacity and Specific Heat
Heat Capacity (Cp) - the amount of heat needed to increase the temperature of an entire object exactly 1 oC
Cop = Joules/oC
Heat capacity is dependent on both the object’s mass and its chemical composition

13. Heat Capacity and Specific Heat
Specific Heat Capacity (c) - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 oC
Often called simply “Specific Heat”
c = Joules/ gram oC

14. Table of Specific Heats

15. Heat Capacity and Specific Heat
Water has a HUGE specific heat value, when is compared to other chemicals:
C (H2O) = 4.18 J/(g oC), or
C (H2O) = 1.00 cal/(g oC)
It heats slowly and
It cools slowly
Consequently, water maintains a relatively constant temperature!

16. Heat Capacity and Specific Heat
Heat calculations for various materials are done with the formula:
 q = mass (g) x c x T
 q = change in heat
c = specific heat
T = change in temperature
Units are either J/(g oC) or cal/(g oC)

17. - Page 299

18. Section 11.2 Measuring and Expressing Enthalpy Changes
OBJECTIVES:
Describe how calorimeters are used to measure heat flow.
Construct thermochemical equations.
Solve for enthalpy changes in chemical reactions by using heats of reaction.

19. Calorimetry
Calorimetry - the precise measurement of the heat into or out of a system for chemical and physical processes.
The device used to measure the absorption or release of heat in chemical or physical processes is called a Calorimeter

20. Calorimetry
Foam cups are excellent heat insulators, and are commonly used as simple calorimeters
For systems at constant pressure, the heat content is the same as a property called Enthalpy (H) of the system

21. A foam cup calorimeter – here, two cups are nestled together for better insulation

22. Calorimetry Changes in heat = Changes in enthalpy = H
q = H These terms are often used interchangeably
Thus, q = H = m x c x T
H is negative for an exothermic process
H is positive for an endothermic process

23. Calorimetry
Calorimetry experiments can be performed at a constant volume using a device called a “bomb calorimeter” - a closed system
Used by nutritionists to measure energy content of food

24. A bomb calorimeter
A Bomb Calorimeter

25. C + O2 → CO2 + 395 kJ ∆ H= -395 kJ C + O2 Energy 395kJ CO2 ® Reactants
Products
C + O2
395kJ
CO2
∆ H= kJ

26. Exothermic
In an exothermic rxn the products are lower in energy than the reactants and energy is released.
ΔHrxn = -395 kJ
The negative sign indicates that energy is released.

27. CaCO3 → CaO + CO2 CaCO3 + 176 kJ → CaO + CO2 CaO + CO2 Energy CaCO3 ®
Reactants
Products
CaO + CO2
176 kJ
CaCO3
∆H = kJ

28. Endothermic
The products are higher in energy than the reactants and energy is absorbed.
ΔHrxn = +176 kJ
The positive sign indicates that energy is absorbed

29. Heat of Reaction
A heat of reaction is the heat change for the equation.
The physical state of reactants and products must also be given.
Standard conditions for the reaction is kPa (1 atm.) and 25 oC (different from STP)

30. Thermochemical Equations
An equation that includes energy is called a thermochemical equation
CH4 + 2O2 ® CO2 + 2H2O kJ
1 mole of CH4 releases kJ of energy.
When you make kJ you also make 2 moles of water

31. CH4 + 2 O2 ® CO2 + 2 H2O kJ
If grams of CH4 are burned completely, how much heat will be produced?
1 mol CH4
802.2 kJ
10. 3 g CH4
16.05 g CH4
1 mol CH4
= 514 kJ
Thus, ΔH = kJ for this reaction

32. Summary, so far...

33. Enthalpy
Enthalpy: The heat content a substance has at a given temperature and pressure
Can’t be measured directly because there is no set starting point
The reactants start with a heat content
The products end up with a heat content
So we can measure how much enthalpy changes
DH (delta H) is the measure of change in heat

34. Enthalpy Enthalpy: Amount of heat in a system (H) .
If heat is released, DH is negative (exothermic)
If heat is absorbed, DH is positive (endothermic)

35. Energy
Change is down
ΔH is <0
Exothermic
Reactants
Products

36. Energy
Change is up
ΔH is > 0
Endothermic
Reactants
Products

37. Heat of Reaction
Heat of Reaction: The heat that is released or absorbed in a chemical reaction
Equivalent to DH
C + O2(g) ® CO2(g) kJ
C + O2(g) ® CO2(g)
DH = kJ
In a thermochemical equation, it is important to indicate the physical state
H2(g) + 1/2O2 (g)® H2O(g) DH = kJ
H2(g) + 1/2O2 (g)® H2O(l) DH = kJ

38. Heat of Combustion C + O2(g) ® CO2(g) + 393.5 kJ C + O2(g) ® CO2(g)
The heat from the reaction that completely burns 1 mole of a substance:
C + O2(g) ® CO2(g) kJ
C + O2(g) ® CO2(g)
DH = kJ

39. Section 17.3 Heat in Changes of State
OBJECTIVES:
Classify the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.

40. Section 17.3 Heat in Changes of State
OBJECTIVES:
Solve for the enthalpy change that occurs when a substance melts, freezes, boils condenses or dissolves.

41. Heat in Changes of State
Heat of Fusion (Hfus) : the heat absorbed by a substance in melting from a solid to a liquid
For H2O: Hfus = 334 J/ g
Molar Hfus = 6.01 kJ/ mol
q = mass x Hfus (there is no temperature change)

42. Molar Heat of Solidification (Hsolid) = heat lost when one mole of liquid solidifies (or freezes) to a solid
For H2O: Hfus = -334 J/ g
Molar Hfus = kJ/ mol
q = mass x Hsolid (no temperature change)

43. Heat in Changes of State
Heat absorbed by a melting solid is equal to heat lost when a liquid solidifies
Thus, Hfus = -Hsolid
Note Table 17.3, page 522

44. - Page 521

45. Heats of Vaporization and Condensation
When liquids absorb heat at their boiling points, they become vapors.
Molar Heat of Vaporization (Hvap) = the amount of heat necessary to vaporize one mole of a given liquid.
q = mass x Hvap (no temperature change)
For H2O: Hfus = 2260 J/ g
Molar Hfus = 40.7 kJ/ mol

46. Heats of Vaporization and Condensation
Condensation is the opposite of vaporization.
Molar Heat of Condensation (Hcond) = amount of heat released when one mole of vapor condenses to a liquid
Hvap = - Hcond

47. Heats of Vaporization and Condensation
Note Figure 17.10, page 523
The large values for water Hvap and Hcond are the reason hot vapors such as steam is very dangerous
You can receive a scalding burn from steam when the heat of condensation is released!
H20(g)  H20(l) Hcond = kJ/mol

48. -Page 524

49. Heat of Solution (Ch 11)
Heat changes can also occur when a solute dissolves in a solvent.
Molar Heat of Solution (Hsoln) = heat change caused by dissolution of one mole of substance
Sodium hydroxide provides a good example of an exothermic molar heat of solution:

50. Heat of Solution NaOH(s)  Na1+(aq) + OH1-(aq) Hsoln = - 445.1 kJ/mol
The heat is released as the ions separate (by dissolving) and interact with water, releasing kJ of heat as Hsoln -thus becoming so hot it steams!
H2O(l)

51. - Page 313

52. End of Chapter 17