1. Chemical Thermodynamics
States of Matter, Phase Changes, Heat, State Functions,
Specific Heats
and
Latent Heats

2. Standards
7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for understanding this concept: a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or atoms). 7. b. Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal energy. 7. c. Students know energy is released when a material condenses or freezes and is absorbed when a material evaporates or melts. 7. d. Students know how to solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change. 7. e.* Students know how to apply Hess’s law to calculate enthalpy change in a reaction. aA + bB = c C + d D ΔHor = [c Hof (C) + dHof (D)] − [a Hof (A) + bHof (B)] 7. f.* Students know how to use the Gibbs free energy equation to determine whether a reaction would be spontaneous. ΔG = Δ H − ΤΔS

5. States of Matter Plasma aka phases Gas Liquid Solid Deionization
Condensation
Boiling
Liquid
Sublimation
Deposition
Freezing
Melting
Solid

6. Condensation (Gas  Liquid)

7. Boiling (Liquid  Gas)

8. Sublimation (Solid  Gas)

9. Deposition (Gas  Solid)

10. Freezing (Liquid  Solid)

11. Melting (Solid  Liquid)

12. Molecular Motion
The state of matter depends on how much energy (motion) the molecules, atoms, or ions have.
The state of matter also depends on how attracted the atoms, molecules, or ions are to each other.

13. Molecular Motion of Gases

14. Molecular Motion
State of Matter O
Gas O H + –
Liquid Cl – Na +
Solid

15. Nonpolar molecules O O

16. Polar molecules O H + – O H + – O H + –

17. Cl – Cl – Na + Na + Cl – Cl – Na + Na + Cl – Cl – Na + Na +
Ionic compounds Cl – Cl – Na + Na + Cl – Cl – Na + Na + Cl – Cl – Na + Na +

18. Temperature Scales Water Boils Human Body Water Freezes 212°F 98.7°F
373K
310K
273K

19. Temperature Scales Surface of Sun Room Temp. Absolute Zero 9,941°F
5,778K
294K 0K

20. Converting Temperatures
Fahrenheit  Celsius
°C = (°F – 32)×(5/9)
Celsius  Fahrenheit
°F = °C ×(9/5) + 32
Celsius  Kelvin
K = °C

21. Absolute Zero
At Zero Kelvin (0 K or – °C), atoms and molecules stop moving.
There is no temperature lower than absolute zero (0 K).

22. Heat – it’s what makes things Hot
Heat is the energy that raises a substance’s temperature.
Heat is transferred from one place to another.
Property
Symbol
Unit
Unit Symbol
Heat Q
calories
cal
Joules J
1 calorie = Joules

23. Heating Water Temperature (°C) Heat Added, Q, (Joules) 100
steam warming
water boiling
water warming
ice melting
ice warming

24. Specific Heat (aka Heat Capacity)
When you heat a substance at a constant rate, its temperature usually increases at a constant rate also.
Q = m·Cp·(ΔT)
heat =
mass ×
specific heat
change in temperature ×
J =
g × °C × J
g °C
Cool Substance +
Warmer Substance
heat

25. ΔT = Tfinal – Tinitial Change in Temperature ΔT = 15˚C – 25˚C
Example: A substance starts with a temperature of 25˚C and ends with a temperature of 15˚C. What is the change in temperature?
ΔT = 15˚C – 25˚C
ΔT = -10˚C

26. Specific Heat Example #1
What is the amount of heat needed to increase the temperature of 4 grams of ethanol by +6˚C?
The specific heat of ethanol is 2.44 J/(g ·˚C). ΔT m Cp
Q = m · Cp · ΔT
Q = (4) · (2.44) · (+6)
Q = 58.6 Joules

27. Specific Heat Example #2Q m
45 Joules of heat are removed from a 0.5 g sheet of aluminum foil that was initially at 120˚C. What will the final temperature of the foil be?
The specific heat of aluminum is 0.90 J/(g ·˚C).
Tinitial Cp
Q = m · Cp · ΔT
ΔT = Tfinal – Tinitial
-45 = (0.5) · (0.90) · ΔT
-45 = 0.45 · ΔT
-100 = Tfinal – 120
ΔT = -100˚C
Tfinal = 20˚C

28. Enthalpy ΔH = enthalpy ΔH = + ΔH = –
Enthalpy – A measure of the total internal heat of a substance per mole.
ΔH = enthalpy
Positive Change in Enthalpy – the process is absorbing heat.
ΔH = +
Negative Change in Enthalpy – the process is releasing heat.
ΔH = –

29. Specific Types of Enthalpy Changes
Heat of Fusion – the heat per mole required to melt a solid.
ΔHfus
Heat of a Reaction – the heat per mole released or absorbed by a reaction.
ΔHr= ΔHproducts – ΔHreactants

30. Exo/Endothermic Reactions
exothermic reaction – the reaction releases heat. (includes freezing, condensation, combustion)
ΔHr = –
Reactants
Products +
heat
endothermic reaction – the reaction absorbs heat. (includes melting, boiling, charging up a battery)
Reactants +
Products
heat
Δ Hr = +

31. N2 (g) + 3 H2 (g) 2 NH3 (g) + 92kJ
nitrogen hydrogen ammonia heat gas gas gas + +

32. Q = m·ΔHfus Q = m·ΔHvap Latent Heat
means hidden
Latent Heat – the heat required to cause a phase-change (hidden because this heat does not change the temperature).
Latent Heat of Fusion – the heat required to melt a solid.
Q = m·ΔHfus
Latent Heat of Vaporization – the heat required to boil a liquid.
Q = m·ΔHvap

33. Heat of Solidification
Latent Heats
Gas / Vapor
Heat of Vaporization
ΔHvap
Heat of Condensation
ΔHcond = – ΔHvap
Liquid
Heat of Fusion
ΔHfus
Heat of Solidification
ΔHsolid = – ΔHfus
Solid

34. Heating Water Temperature (°C) Heat Added, Q, (Joules) 100
Q = m·(ΔT)·Cp steam
steam warming
Q = m·ΔHvap
water boiling
Q = m·(ΔT)·Cp water
water warming
Q = m·ΔHfus
ice melting
Q = m·Cp ice·(ΔT)
ice warming

35. Using Heats of Formation
Heat of Formation (ΔH0f) – A measure of the internal heat of a substance relative to other substances.
compared to
As a reference point, the heat of formation of an element (ex. Fe, O2, H2) is 0 kJ/mol.
ΔHr= ΔH0f products – ΔH0f reactants

36. ΔHr= - (B.E. products – B.E. reactants)
Using Bond Energies
Bond Energy (B.E.) – A measure of energy that is locked away in covalent bonds.
There are different bond energies depending on the types of atoms and if it is a single bond, double bond, etc.
ΔHr= - (B.E. products – B.E. reactants)
because we are breaking bonds

37. Lewis Dot Structures O H C O C O N O O C H

38. Entropy ΔS = entropy ΔS = + ΔS = –
Entropy – A measure of the disorder of a substance.
ΔS = entropy
Positive Change in Entropy – the substances are becoming more disordered.
ΔS = +
Negative Change in Entropy – the substances are becoming less disordered.
ΔS = –
Nature tends to favor a continual increase in entropy unless outside energy is used.
(ex. Cleaning a room, photosynthesis, distillation)

39. Which has the greater entropy?
Low Entropy
High Entropy

40. Which has the greater entropy?
High Entropy
Low Entropy

41. Which has the greater entropy?
High Entropy
Low Entropy

42. Which has the greater entropy?
High Entropy
Low Entropy

43. Which has the greater entropy?
Low Entropy
High Entropy

44. Gibbs (Free) Energy ΔG = Gibbs energy ΔG = ΔH – ΔS × T ΔG = – ΔG = +
Gibbs Energy – A measure of both the enthalpy and entropy of a substance combined.
ΔG = Gibbs energy
ΔG = ΔH – ΔS × T
Negative Change in Gibbs energy – a process can happen. The process is spontaneous, but the process might still be incredibly slow.
ΔG = –
Positive Change in Gibbs energy – a process will not happen. The process is nonspontaneous.
ΔG = +

45. State Functions
Enthalpy (ΔH), Entropy (ΔS), and Gibbs Energy (ΔG) are all state functions.
They are called state functions because it doesn’t matter how they got there, it only matters where they started and where they finished.

46. State Functions X X

47. X X State Functions Water at 100˚C Ice at -30˚C Water at 20˚C

48. State Functions Summary
Property
Symbol
Meaning
Units
Sign change
Enthalpy ΔH
heat
+ = endothermic
– = exothermic
Entropy ΔS
disorder
+ = increase in disorder
– = decrease in disorder
Gibbs energy ΔG
does it happen?
+ = nonspontaneous
– = spontaneous
0 = at equilibrium kJ
mol kJ
mol·K kJ
mol

49. Exo/Endothermic Reactions
exothermic reaction – the reaction releases heat.
endothermic reaction – the reaction absorbs heat.
H2O (g) + 3 H2 (g) 2 NH3 (g) + 92kJ
Solid hydrogen ammonia heat gas gas gas

51. He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe

52. He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe

54. 4 e– in valence shell