1. Chapter 19 Chemical Thermodynamics
CHEMISTRY The Central Science 9th Edition
Chapter 19 Chemical Thermodynamics
David P. White
Prentice Hall © 2003
Chapter 19

2. Spontaneous Processes
Thermodynamics is concerned with the question: can a reaction occur?
First Law of Thermodynamics: energy is conserved.
Any process that occurs without outside intervention is spontaneous.
When two eggs are dropped they spontaneously break.
The reverse reaction is not spontaneous.
A process that is spontaneous in one direction is not spontaneous in the opposite direction.
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3. Spontaneity & Temperature
The direction of a spontaneous process can depend on temperature: ice turning to water is spontaneous at T > 0C, water turning to ice is spontaneous at T < 0C.

4. Packet Example Page 3
Predict whether the following processes are spontaneous as described, spontaneous in the reverse direction, or are in equilibrium.
a) When a piece of metal heated to 150C is added to water at 40C, the water gets hotter.
b) Water at room temperature decomposed into hydrogen and oxygen gases.
c) Benzene vapor at a pressure of 1 atm condenses to liquid benzene at the normal boiling point of benzene.
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5. Reversible Processes
Reversible processes can go back and forth between states along the same path.
Example: changes of state
Path taken back to original state is exactly the reverse of the forward process.
No net change in system or surroundings when cycle is completed.
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6. Reversible Processes
When 1 mol of water is frozen at 1 atm at 0C to form 1 mol of ice, q = Hvap of heat is removed.
To reverse the process, q = Hvap must be added to the 1 mol of ice at 0C and 1 atm to form 1 mol of water at 0C.
Converting between 1 mol of ice and 1 mol of water at 0C is a reversible process.

7. Irreversible Processes
An irreversible process cannot be reversed to restore the system and surroundings back to their original state.
A different path with different values of q and w are needed. Surroundings are not returned to original conditions.

8. Reversible & Irreversible Processes
Chemical systems in equilibrium are reversible.
In any spontaneous process, the path between reactants and products is irreversible.
Thermodynamics gives us the direction of a process but it cannot predict the speed at which the process will occur.
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9. Spontaneous Expansion of a Gas
Why does the gas expand?
Why is the reverse process nonspontaneous?

10. Spontaneous Expansion of a Gas
Consider 2 gas molecules in one flask (a).
Once the stopcock is open, there is a higher probability that one molecule will be in each flask than both molecules being in the same flask (b).

11. Entropy Entropy, S, is a measure of the disorder of a system.
The more disordered or random the system, the large value of S.
Spontaneous reactions proceed to lower energy or higher entropy.
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12. In ice, molecules are very well ordered because of the H-bonds.
Ice has a low entropy.
As ice melts, intermolecular forces are broken , order is interrupted.
Water is more random than ice, has higher entropy.
Ice spontaneously melts at room temperature.

13. There is a balance between energy and entropy considerations.
When an ionic solid is placed in water two things happen:
the water organizes into hydrates about the ions (so the entropy decreases), and
the ions in the crystal dissociate (the hydrated ions are less ordered than the crystal, so the entropy increases).
A balance between energy and entropy considerations.
When an ionic solid is placed in water two things happen:
Hydrates form (entropy decreases)
Ions dissociate (entropy increases)

14. Entropy, Cont.
Generally, when an increase in entropy in one process is associated with a decrease in entropy in another, the increase in entropy dominates.
Entropy is a state function.
For a system, S = Sfinal - Sinitial.
If S > 0 the randomness increases, if S < 0 the order increases.
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15. Packet Example Page 6
By considering the disorder in the reactants and products, predict whether ∆S is positive or negative for the following:
a) H2O (l) → H2O (g)
b) Ag+(aq) + Cl-(aq) → AgCl (s)
c) 4Fe(s) + 3O2 (g) → 2Fe2O3 (s)
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16. Second Law of Thermodynamics
Second Law of Thermodynamics: In any spontaneous process, the entropy of the universe increases.
Suniv = Ssys + Ssurr
Entropy is not conserved: Suniv is increasing.
To predict spontaneity, we must know the sign of ∆Suniv.
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17. Entropy and the Second Law of Thermodynamics
For a reversible process: Suniv = 0.
For a spontaneous process (i.e. irreversible): Suniv > 0. Suniv< 0 means spontaneous is opposite direction.
Second law states that the entropy of the universe must increase in a spontaneous process. Entropy of a system can decrease as long as the entropy of the surroundings increases.
For an isolated system, Ssys = 0 for a reversible process and Ssys > 0 for a spontaneous process.

18. Entropy and the Second Law of Thermodynamics
Suppose a system changes reversibly between state 1 and state 2. At constant T where qrev is the amount of heat added reversibly to the system, the change in entropy is given by
Example: a phase change occurs at constant T with the reversible addition of heat.
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19. Example Packet p 7
The element mercury, Hg, is a silvery liquid at room temperature. The normal freezing point of mercury is -38.9C and its molar enthalpy of fusion is ∆H = kJ/mol. Calculate the entropy change of the system when 50.0 g liquid mercury freezes at the normal freezing point.
Remember: fusion = melting
Use:

20. Practice Example
The normal boiling point of ethanol (C2H5OH) is 78.3C and its ∆H vaporization= kJ/mol. Calculate ∆S when 68.3 g of ethanol at 1 atm condense to liquid at the normal boiling point.
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21. Example Packet p 9
Consider the reversible melting of 1 mol of ice in a large, isothermal water bath at 0C. The enthalpy of fusion is 6.01 kJ/mol. Calculate the entropy change in the system and in the surroundings and the overall change in entropy of the universe for this process.
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22. Practice Example
∆Hvap = kJ/mol for liquid bromine. Calculate the change in entropy for the system, the surroundings, and the universe for the reversible vaporization of liquid bromine at its normal boiling point of 59C.
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23. Effect of Temperature on Spontaneity
∆Ssurr are determined primarily by flow of energy into and out of system as heat.
An exothermic process in system increases entropy in surroundings.
Exothermicity is a driving force for spontaneity BUT it depends on the temperature at which the process occurs!
Impact of transfer of a given quantity of energy as heat to or from surroundings will be greater at lower temperatures.
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24. Entropy Changes in the Surroundings
Sign of ∆Ssurr depends on direction of heat flow.
Magnitude of ∆Ssurr depends on the temperature.
∆Ssurr depends directly on quantity of heat transferred and inversely on temperature.
∆Ssurr = -∆H / T
(We need a minus sign here to change the point of view from the system to the surroundings.)
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25. Examples from Zumdahl pp 782 & 783
Calculate ∆Ssurr for the following reactions at 25C and
1 atm.
Sb2S3(s) + 3Fe(s) → 2Sb(s) + 3FeS(s) ∆H = -125 kJ
Sb4O6(s) + 6C(s) → 4Sb(s) + 6CO(s) ∆H = +778 kJ
Use ∆Ssurr = -∆H / T
+419 J/K
-2.61 x 103 J/K
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26. Determining the Sign of ∆Suniv
Signs of Entropy Changes
∆Ssys ∆Ssurr ∆Suniv Process Spontaneous?
Yes
No, rxn goes in reverse
? Yes, if ∆Ssys  ∆Ssurr
? Yes, if ∆Ssurr  ∆Ssys
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27. Atomic Modes of Motion
Movement of molecules is related to energy and entropy.
Three atomic modes of motion:
translation (moving from one point in space to another),
vibration (shortening and lengthening of bonds, including the change in bond angles),
rotation (spinning around an axis).

28. The Molecular Interpretation of Entropy
Energy is required to get a molecule to translate, vibrate or rotate.
The more energy stored in translation, vibration and rotation, the greater the degrees of freedom and the higher the entropy.
In a perfect crystal at 0 K there is no translation, rotation or vibration of molecules. Therefore, this is a state of perfect order.
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29. Third Law of Thermodynamics
Third Law of Thermodynamics: the entropy of a perfect crystal at 0 K is zero.
As we heat a substance from absolute zero, the entropy must increase.
Entropy changes dramatically at a phase change.
Entropy increases when
--Liquids or solutions are formed from solids
--Gases are formed from solids or liquids
--The number of gas molecules increases.

30. Heat and Entropy
As we heat a substance from absolute zero, the entropy must increase.

31. Boiling corresponds to a much greater change in entropy than melting.

32. Example Packet p 11
Choose the sample of matter that has the greater entropy in each pair and explain your choice:
a) 1 mol of solid NaCl or 1 mol of gaseous HCl at 25C
b) 2 mol of HCl(g) or 1 mol of HCl(g) at 25C
c) 1 mol of HCl(g) or 1 mol of Ar(g) at 25C
d) 1 mol of N2(s) at 24 K or 1 mol of N2 at 25C
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33. Example Packet p 12
Predict whether the entropy change of the system in each of the following isothermal reactions is positive or negative.
a) CaCO3(s) → CaO(s) + CO2(g)
b) N2(g) + 3H2(g) → 2NH3(g)
c) N2(g) + O2(g) → 2NO(g)
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34. Entropy Changes in Chemical Reactions
Absolute entropy can be determined from complicated measurements.
Standard molar entropy, S: entropy of a substance in its standard state. Similar in concept to H.
Units: J/mol-K. Note units of H: kJ/mol.
Standard molar entropies of elements are not zero.
For a chemical reaction which produces n moles of products from m moles of reactants:

36. Gibbs Free Energy
For a spontaneous reaction the entropy of the universe must increase.
Reactions with large negative H values are spontaneous.
How to we balance S and H to predict whether a reaction is spontaneous?
Gibbs free energy, G, of a state is
For a process occurring at constant temperature
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37. Gibbs Free Energy There are three important conditions:
If G < 0 then the forward reaction is spontaneous.
If G = 0 then reaction is at equilibrium and no net reaction will occur.
If G > 0 then the forward reaction is not spontaneous. If G > 0, work must be supplied from the surroundings to drive the reaction.
For a reaction the free energy of the reactants decreases to a minimum (equilibrium) and then increases to the free energy of the products.
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38. Gibbs Free Energy
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39. Gibbs Free Energy
Consider the formation of ammonia from nitrogen and hydrogen:
Initially ammonia will be produced spontaneously (Q < Keq).
After some time, the ammonia will spontaneously react to form N2 and H2 (Q > Keq).
At equilibrium, ∆G = 0 and Q = Keq.
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41. Gibbs Free Energy Standard Free-Energy Changes
We can tabulate standard free-energies of formation, Gf (c.f. standard enthalpies of formation).
Standard states are: pure solid, pure liquid, 1 atm (gas), 1 M concentration (solution), and G = 0 for elements.
G for a process is given by
The quantity G for a reaction tells us whether a mixture of substances will spontaneously react to produce more reactants (G > 0) or products (G < 0).
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42. Free Energy and Temperature
Focus on G = H - TS:
If H < 0 and S > 0, then G is always negative.
If H > 0 and S < 0, then G is always positive. (That is, the reverse of 1.)
If H < 0 and S < 0, then G is negative at low temperatures.
If H > 0 and S > 0, then G is negative at high temperatures.
Even though a reaction has a negative G it may occur too slowly to be observed.
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43. Free Energy and Temperature
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44. Free Energy and The Equilibrium Constant
Recall that G and K (equilibrium constant) apply to standard conditions.
Recall that G and Q (equilibrium quotient) apply to any conditions.
It is useful to determine whether substances under any conditions will react:
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45. Free Energy and The Equilibrium Constant
At equilibrium, Q = K and G = 0, so
From the above we can conclude:
If G < 0, then K > 1.
If G = 0, then K = 1.
If G > 0, then K < 1.
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46. End of Chapter 19 Chemical Thermodynamics
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