1. Thermodynamics Standard 7
Chemistry.
Ms. Siddall.

2. Chemical Thermodynamics = the movement of heat in a chemical reaction.
Standard 7a: ‘heat flow’
Chemical Thermodynamics = the movement of heat in a chemical reaction.
Temperature = a measure of the average kinetic energy of particle motion
Heat = The transfer of energy from a hotter object to a colder object (sometimes called ‘heat flow’)
temperature measures energy
Heat measures energy transfer

3. Summary 1
Describe the difference between heat and temperature

4. Energy transfer Summary 2
Particle vibrations increase when a particle gains energy
Vibrations are transferred to surrounding particles
Summary 2
Describe how energy is transferred between atoms.

5. Identifying heat transfer: System: experiences a change
Surroundings: causes a change
e.x. hot coffee (system) cools because it transfers heat to the air, the cup, the table & the whole universe! (surroundings)

6. Summary 3 Consider an ice cube dropped into a glass of warm water.
Ice cube = system
Water = surroundings
Does heat flow into the system or out of the system?
What is gaining energy (system or surroundings)?

7. Summary 4 Endothermic Process: A process in which energy is absorbed.
Standard 7b: exothermic & endothermic process
Endothermic Process: A process in which energy is absorbed.
Example: Water boiling
H2O(l) + heat  H2O(g)
In an endothermic process heat is a reactant.
reactants
product
Summary 4
In an endothermic process which has more energy; reactants or products?

8. Summary 5 Exothermic Process: A process in which energy is released.
Example:
A fire
3C + 2O2  heat + 2CO + CO2
In an exothermic process heat is a product
products
reactants
Summary 5
In an exothermic process which has more energy, reactants or products?

9. Energy diagram Summary 6 H2O(g) Increasing energy H2O(l)
Draw an energy diagram for the campfire reaction.
Show reactants and products.
Draw only one arrow from reactants to products and label the arrow (endothermic or exothermic)
H2O(g)
exothermic
endothermic
Increasing energy
H2O(l)

10. Transition State energy diagram
activation energy = energy needed to form transition state (activated complex)
Transition state
Energy released when products form
reactants
Total energy released during reaction
energy
products

11. Transition State: An intermediate state that can occur during a reaction
Also called an ‘activated complex’
An exothermic reaction is not always spontaneous because energy is needed to form a transition state.
e.x. a spark is needed to start a fire
Summary 7
Draw a transition state energy diagram for an endothermic reaction

12. Measuring heat flow.
Energy is measured in joules (J) or calories (cal)
Example: 334J of energy are needed to melt 1g of ice.
1 calorie (c) = 4.18J
1 food calorie (C) = 1000 calories = 4180J
Summary 8
If your body burns about 2,000 food calories a day, approximately how many joules of energy is that?

13. Energy released = exothermic
KJ = kilojoules = 1000J
Showing a change in energy:
S(s) + O2(g)  SO2(g) + energy
S(s) + O2(g)  SO2(g) + 297KJ
S(s) + O2(g)  SO2(g) ∆H = -297KJ
-∆H = exothermic
+∆H = endothermic
∆H = change in enthalpy
Enthalpy = energy/heat

14. N2(g) + 2O2(g)  2NO2(g) ∆H = + 68KJ
N2(g) + 2O2(g) + 68KJ  2NO2(g)
Endothermic reaction
Energy is a reactant

15. Summary 9
Write an equation to show water melting. Use ∆H to show energy.
(it takes 5.9kJ of energy to melt ice)
Is ∆H negative or positive? Why?

16. Phase Change: The physical state of a compound changes
Standard 7c: energy of phase change
Phase Change: The physical state of a compound changes
The same compound is observed before and after the change
Example: ice melting H2O(s)  H20(l)
There is no temperature change.
Energy is used to overcome intermolecular attractions.

17. Summary 10
Is the example of ice melting an endothermic process or an exothermic process?

18. gas liquid solid Physical state endothermic exothermic evaporating
Condensing
break hydrogen bonds
endothermic
liquid
exothermic
Energy released intermolecular attractions take over
melting
freezing
break lattice structure
solid

19. Summary 11
In which phase do the molecules have the most energy? (solid, liquid, or gas)
Is the process of condensing endothermic or exothermic?
Is the process of vaporization endothermic or exothermic?

20. Freezing/boiling point graph for water.
Standard 7d: solving problems
Energy absorbed
= no temp change
= physical change
boiling
ΔHvap
110
100
steam
melting ΔHfus
Water (CH2O(l))
Temperature (°C)
Energy absorbed
= Change in temperature
= Change in K.E.
ice
-10
energy

21. Summary 12 Which two sections of the graph show no temperature change.
Why is there no temperature change in these sections?

22. ΔHfus = enthalpy of fusion (J/g)
Standard 7d: solving problems
Latent Heat of fusion. (latent heat = hidden heat)
ΔHfus = The energy released when 1g of a substance is frozen OR the energy needed when 1g of a substance is melted.
ΔHfus = enthalpy of fusion (J/g)
Fusion = freezing (liquid  solid)
also used for melting (solid  liquid)
Summary 13: What does ‘fusion’ mean?

23. Example: freezing water
How much energy is released when 10g water freezes? (ΔHfusH2O = 334J/g)
10g H2O(s)
334J
= 3340J
= 3.34kJ
1g H2O(s)
Summary 14
How much energy is needed to melt 100g of water? (show calculation)

24. ΔHvap = enthalpy of vaporization (J/g)
Latent Heat of vaporization
ΔHvap = The energy needed when 1g of a substance is evaporated OR the energy released when 1g of a substance is condensed.
ΔHvap = enthalpy of vaporization (J/g)
vaporization = evaporating (liquid  gas)
also used for condensing (gas liquid)

25. Summary 15
What does vaporization mean?
What does condensation mean?

26. Summary 16 Example: Boiling water
How much energy is needed to boil 10g water? (ΔHvapH2O = 2260J/g)
10g H2O(l)
2260J
= 22600J
= 22.6kJ
= J
1g H2O(l)
Summary 16
How much energy is released when 100g of water vapor is condensed? (show work)

27. Heat Capacity. C = specific heat capacity Example: CH2O(l) = 4.18J/g°C
The amount of heat energy needed to raise the temperature of 1g of a substance by 1°C
Example: CH2O(l) = 4.18J/g°C
It takes 4.18J of energy to raise the temperature of 1g of water by 1°C
1 calorie = 4.18J

28. Summary 17
How much energy is needed to raise the temperature of 1g of water by 1°C? (give your answer in joules and calories)

29. Example.
How much energy is needed to raise the temperature of 5g water from 22°C to 24°C? (CH2O(l) = 4.18J/g°C)
5g H2O(l)
4.18J H2O(l)
2°C
= 41.8J
= J
g °C

30. Summary 18
How much energy is released when 10g water cools from 40°C to 30°C?
10g H2O(l)
4.18J H2O(l)
10°C
= 418.J
= J
g °C

31. Measuring specific heat capacity for different compounds
Thermometer: Measures temperature change for water
‘q’=energy released by metal =energy absorbed by water
Unknown compound: heated to 100°C and placed in the cold water
H2O

32. Summary 19
How much energy (q) is released by a metal if the temperature of 100g of water in the calorimeter rises from 20°C to 30°C?

33. Measuring the heat of a reaction: ‘q’
(q = energy released or absorbed by water)
Thermometer: measures temperature change for water
T  = exothermic
T  = endothermic
Reaction chamber:
3H2 + N2  NH3
heat of reaction is absorbed by water
100g H2O

34. Example: 10g NH3 are produced in the above reaction
Example: 10g NH3 are produced in the above reaction. The temperature rises from 20.0°C to 30.0°C.
Calculate ‘q’ (energy) for the reaction.
Is the reaction endothermic or exothermic?
Calculate ΔH (J/g) for this reaction
Calculate ΔH (mol/g) for this reaction

35. Kinetic energy distribution diagram

36. Kinetic energy distribution diagram
T1 = low temperature = low energy
T2 = higher temperature = higher energy
Emin = minimum energy needed to escape.
More T2 particles have Emin
Less T1 particles have Emin

37. Summary 20
Explain why more particles evaporate from a cup of hot water compared to a cup of cold water.

38. Find the sum of the 2 equations…
Standard 7e: Apply Hess’s Law to calculate enthalpy change in a reaction
Hess’s Law: If a series of reactions are added together the enthalpy change for the net reaction will be the sum of the enthalpy changes for the individual steps.
E.x. N2(g) + 2O2(g)  2NO2(g)
N2(g) + O2(g)  2NO(g) ΔH = +181kJ
2NO(g) + O2(g)  2NO2(g) ΔH = -113kJ
Find the sum of the 2 equations…

39. N2(g) + O2(g)  2NO(g) ΔH = +181kJ
2NO(g) + O2(g)  2NO2(g) ΔH = -113kJ
N2(g) + O2(g) + 2NO(g) + O2(g)  2NO(g) + 2NO2(g)
N2(g) + 2O2(g)  2NO2(g)
ΔH =
notes:
You can reverse reactions (change sign of ΔH)
You can multiply or divide equations (do same to ΔH)
+181kJ +
(-113kJ) =
+68kJ

40. Hess summary
Complete questions 66, 74, 81 & 84 on page 536 & 537