1. Section 11.1 The Flow of Energy - Heat
OBJECTIVES:
Explain the relationship between energy and heat.

2. Section 11.1 The Flow of Energy - Heat
OBJECTIVES:
Distinguish between heat capacity and specific heat.

3. Energy and Heat
Thermochemistry - concerned with heat changes that occur during chemical reactions
Energy - capacity for doing work or supplying heat
weightless, odorless, tasteless
if within the chemical substances- called chemical potential energy

4. Energy and Heat
Gasoline contains a significant amount of chemical potential energy
Heat - represented by “q”, is energy that transfers from one object to another, because of a temperature difference between them.
only changes can be detected!
flows from warmer  cooler object

5. Exothermic and Endothermic Processes
Essentially all chemical reactions, and changes in physical state, involve either:
release of heat, or
absorption of heat

6. Exothermic and Endothermic Processes
In studying heat changes, think of defining these two parts:
the system - the part of the universe on which you focus your attention
the surroundings - includes everything else in the universe

7. Exothermic and Endothermic Processes
Together, the system and it’s surroundings constitute the universe
Thermochemistry is concerned with the flow of heat from the system to it’s surroundings, and vice-versa.
Figure 11.3, page 294

8. Exothermic and Endothermic Processes
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.

9. Exothermic and Endothermic Processes
Fig. 11.3a, p heat flowing into a system from it’s surroundings:
defined as positive
q has a positive value
called endothermic
system gains heat as the surroundings cool down

10. Exothermic and Endothermic Processes
Fig. 11.3b, p heat flowing out of a system into it’s surroundings:
defined as negative
q has a negative value
called exothermic
system loses heat as the surroundings heat up

11. Exothermic and Endothermic
Fig. 11.4, page on the left, the system (the people) gain heat from it’s surroundings (the fire)
this is endothermic
On the right, the system (the body) cools as perspiration evaporates, and heat flows to the surroundings
this is exothermic

12. Exothemic and Endothermic
Every reaction has an energy change associated with it
Exothermic reactions release energy, usually in the form of heat.
Endothermic reactions absorb energy
Energy is stored in bonds between atoms

13. Heat Capacity and Specific Heat
A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1 oC.
Used except when referring to food
a Calorie, written with a capital C, always refers to the energy in food
1 Calorie = 1 kilocalorie = 1000 cal.

14. Heat Capacity and Specific Heat
The calorie is also related to the joule, the SI unit of heat and energy
named after James Prescott Joule
4.184 J = 1 cal
Heat Capacity - the amount of heat needed to increase the temperature of an object exactly 1 oC

15. Heat Capacity and Specific Heat
Specific Heat Capacity - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 oC (abbreviated “C”)
often called simply “Specific Heat”
Note Table 11.2, page 296
Water has a HUGE value, compared to other chemicals

16. Heat Capacity and Specific Heat
For water, C = 4.18 J/(g oC), and also C = 1.00 cal/(g oC)
Thus, for water:
it takes a long time to heat up, and
it takes a long time to cool off!
Water is used as a coolant!
Note Figure 11.7, page 297

17. Heat Capacity and Specific Heat
To calculate, use the formula:
q = mass (g) x T x C
heat abbreviated as “q”
T = change in temperature
C = Specific Heat
Units are either J/(g oC) or cal/(g oC)
Sample problem 11-1, page 299

18. Section 11.2 Measuring and Expressing Heat Changes
OBJECTIVES:
Construct equations that show the heat changes for chemical and physical processes.

19. Section 11.2 Measuring and Expressing Heat Changes
OBJECTIVES:
Calculate heat changes in chemical and physical processes.

20. Calorimetry
Calorimetry - the accurate and precise measurement of heat change 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

21. Calorimetry
Foam cups are excellent heat insulators, and are commonly used as simple calorimeters
Fig. 11.8, page 300
For systems at constant pressure, the heat content is the same as a property called Enthalpy (H) of the system

22. Calorimetry Changes in enthalpy = H
q = H These terms will be used interchangeably in this textbook
Thus, q = H = m x C x T
H is negative for an exothermic reaction
H is positive for an endothermic reaction (Note Table 11.3, p.301)

23. Calorimetry
Calorimetry experiments can be performed at a constant volume using a device called a “bomb calorimeter” - a closed system
Figure 11.9, page 301
Sample 11-2, page 302

24. C + O2 ® CO2
+ 395 kJ
Energy
Reactants
Products
C + O2
395kJ
C O2

25. In terms of bonds C O O C O Breaking this bond will require energy. C
Making these bonds gives you energy.
In this case making the bonds gives you more energy than breaking them.

26. Exothermic The products are lower in energy than the reactants
Releases energy

27. CaCO3 + 176 kJ ® CaO + CO2 CaCO3 ® CaO + CO2 Energy Reactants Products

28. Endothermic The products are higher in energy than the reactants
Absorbs energy
Note Figure 11.11, page 303

29. Chemistry Happens in MOLES
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

30. Thermochemical Equations
A heat of reaction is the heat change for the equation, exactly as written
The physical state of reactants and products must also be given.
Standard conditions for the reaction is kPa (1 atm.) and 25 oC

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

32. How many grams of water would be produced with 506 kJ of heat? (22.7g)
CH4 + 2 O2 ® CO2 + 2 H2O kJ
How many liters of O2 at STP would be required to produce 23 kJ of heat? (1.28L)
How many grams of water would be produced with 506 kJ of heat? (22.7g)

33. Summary, so far...

34. 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

35. Enthalpy Symbol is H Change in enthalpy is DH (delta H)
If heat is released, the heat content of the products is lower
DH is negative (exothermic)
If heat is absorbed, the heat content of the products is higher
DH is positive (endothermic)

36. Energy
Change is down
DH is <0
Reactants
Products

37. Energy
Change is up
DH is > 0
Reactants
Products

38. 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 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

39. Heat of Combustion
The heat from the reaction that completely burns 1 mole of a substance
Note Table 11.4, page 305

40. Section 11.3 Heat in Changes of State
OBJECTIVES:
Classify, by type, the heat changes that occur during melting, freezing, boiling, and condensing.

41. Section 11.3 Heat in Changes of State
OBJECTIVES:
Calculate heat changes that occur during melting, freezing, boiling, and condensing.

42. Heats of Fusion and Solidification
Heat of Fusion (Hfus) - the heat absorbed by a gram of a substance in melting from a solid to a liquid
Heat of Solidification (Hsolid) - heat lost when one gram of liquid solidifies

43. Heats of Fusion and Solidification
Heat absorbed by a melting solid is equal to heat lost when a liquid solidifies
Thus, Hfus = -Hsolid

44. Heats of Vaporization and Condensation
When liquids absorb heat at their boiling points, they become vapors.
Heat of Vaporization (Hvap) - the amount of heat necessary to vaporize one gram of a given liquid.

45. Heats of Vaporization and Condensation
Condensation is the opposite of vaporization.
Heat of Condensation (Hcond) - amount of heat released when one gram of vapor condenses
Hvap = - Hcond

46. Heats of Vaporization and Condensation
Note Figure 11.5, page 310
The large values for 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!

47. Heats of Vaporization and Condensation
H20(g)  H20(l) Hcond = - 2,260J/g
How much heat is released by 34.0 grams of water vapor when it condenses to liquid water?
(34.0g)(-2,260J/g) = -76,800J

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

49. Heat of Solution NaOH(s)  Na1+(aq) + OH1-(aq) Hsoln = - 44.5 kJ/mol
The heat is released as the ions separate and interact with water, releasing 44.5 kJ of heat as Hsoln thus becoming so hot it steams!
H2O(l)