6.1 What is Energetics?

1
Energetics
6.1 What is Energetics? 6.2 Enthalpy Changes
Related to Breaking and Forming of
Bonds 6.3 Standard Enthalpy Changes 6.4 Experiment
al Determination of Enthalpy Changes by
Calorimetry 6.5 Hesss Law 6.6 Calculations
involving Standard Enthalpy Changes of Reactions
2
What is energetics?
6.1 What is energetics? (SB p.136)
Energetics is the study of energy changes
associated with chemical reactions.
Thermochemistry is the study of heat changes
associated with chemical reactions.
Some terms
Enthalpy(H) heat content in a substance
Enthalpy change(?H) heat content of products
- heat content of reactants Hp - Hr
3
Law of conservation of energy
6.1 What is energetics? (SB p.136)
The law of conservation of energy states that
energy can neither be created nor destroyed.
4
Internal energy and enthalpy
6.1 What is energetics? (SB p.137)
e.g. Zn(s) 2HCl(aq) ? ZnCl2(aq) H2(g)
5
Internal energy and enthalpy
6.1 What is energetics? (SB p.138)
Enthalpy change
(Heat change at constant volume)
6
Exothermic and endothermic reactions
6.1 What is energetics? (SB p.138)
An exothermic reaction is a reaction that
releases heat energy to the surroundings. (?H
-ve)
7
Exothermic and endothermic reactions
6.1 What is energetics? (SB p.139)
An endothermic reaction is a reaction that
absorbs heat energy from the surroundings. (?H
ve)
8
Energy Changes Related to Breaking and Forming of
Bonds
9
Enthalpy changes related to breaking and forming
of bonds
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.140)
e.g. CH4 2O2? CO2 2H2O
10
Enthalpy changes related to breaking and forming
of bonds
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.140)
In an exothermic reaction, the energy required in
breaking the bonds in the reactants is less than
the energy released in forming the bonds in the
products (products contain stronger bonds).
11
Enthalpy changes related to breaking and forming
of bonds
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.140)
12
Enthalpy changes related to breaking and forming
of bonds
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.140)
In an endothermic reaction, the energy required
in breaking the bonds in the reactants is more
than the energy released in forming the bonds in
the products (reactants contain stronger bonds).
13
Standard Enthalpy Changes
14
Standard enthalpy changes
6.3 Standard enthalpy changes (SB p.141)
CH4(g) 2O2(g) ? CO2(g) 2H2O(g) ?H -802
kJ mol-1
CH4(g) 2O2(g) ? CO2(g) 2H2O(l) ?H
-890 kJ mol-1
15
Standard enthalpy changes
6.3 Standard enthalpy changes (SB p.141)
As enthalpy changes depend on temperature and
pressure, it is necessary to define standard
states and conditions
1. elements or compounds in their normal
physical states2. a pressure of 1 atm (101325
Nm-2) and3. a temperature of 25oC (298 K)
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Standard enthalpy changes of neutralization
6.3 Standard enthalpy changes (SB p.142)
17
Standard enthalpy changes of neutralization
6.3 Standard enthalpy changes (SB p.142)
H(aq) OH-(aq) ? H2O(l)
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Standard enthalpy changes of neutralization
6.3 Standard enthalpy changes (SB p.142)
Enthalpy level diagram for the neutralization of
a strong acid and a strong alkali
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Standard enthalpy change of solution
6.3 Standard enthalpy changes (SB p.142)
20
Standard enthalpy change of solution
6.3 Standard enthalpy changes (SB p.143)
21
Standard enthalpy change of solution
6.3 Standard enthalpy changes (SB p.143)
22
Standard enthalpy change of solution
6.3 Standard enthalpy changes (SB p.143)
23
Standard enthalpy change of formation
6.3 Standard enthalpy changes (SB p.144)
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Standard enthalpy change of formation
6.3 Standard enthalpy changes (SB p.144)
25
Standard enthalpy change of formation
6.3 Standard enthalpy changes (SB p.144)
N2(g) ? N2(g)
The enthalpy change of formation of an element is
always zero.
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Standard enthalpy change of combustion
6.3 Standard enthalpy changes (SB p.146)
e.g. C3H8(g) 5O2(g) ? 3CO2(g) 4H2O(l) ?H1
-2220 kJ
2C3H8(g) 10O2(g) ?6CO2(g) 8H2O(l)
?H2 ?
?H2 -4440 kJ
? It is more convenient to report enthalpy
changes per mole of the main reactant
reacted/product formed.
27
Standard enthalpy change of combustion
6.3 Standard enthalpy changes (SB p.146)
28
Standard enthalpy change of combustion
6.3 Standard enthalpy changes (SB p.147)
29
Experimental Determination of Enthalpy Changes by
Calorimetry
30
Experimental determination of enthalpy changes by
calorimetry
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.148)
Calorimeter a container used for measuring the
temperature change of solution
31
Determination of enthalpy change of neutralization
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.149)
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6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.149)
Heat evolved (m1c1 m2c2) ?T where m1 is the
mass of the solution, m2 is the mass of
calorimeter, c1 is the specific heat capacity of
the solution, c2 is the specific heat capacity
of calorimeter, ?T is the temperature change of
the reaction
33
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.150)
Determination of enthalpy change of combustion
The Philip Harris calorimeter used for
determining the enthalpy change of combustion of
a liquid fuel
34
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.151)
Determination of enthalpy change of combustion
A simple apparatus used to determine the enthalpy
change of combustion of ethanol
35
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.151)
Heat evolved (m1c1 m2c2) ?T Where m1 is the
mass of water in the calorimeter, m2 is the
mass of the calorimeter, c1 is the specific heat
capacity of the water, c2 is the specific heat
capacity of calorimeter, ?T is the temperature
change of the reaction
36
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.152)
Determination of enthalpy change of solution

• By measuring the temperature change when a known
  mass of solute is added to a known volume of
  solvent in a calorimeter
• Heat change (m1c1 m2c2) ?T

37
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Determination of enthalpy change of formation

• The enthalpy change of formation of a substance
  can be quite high
• Found out by applying Hesss law of constant heat
  summation

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Hesss Law
39
Hesss Law
6.5 Hesss law (SB p.153)
A B
C D
?H1 ?H2 ?H3
Hesss law of constant heat summation states that
the total enthalpy change accompanying a chemical
reaction is independent of the route by which the
chemical reaction takes place.
40
Enthalpy level diagram
6.5 Hesss law (SB p.154)

• Relate substances together in terms of enthalpy
  changes of reactions

41
Enthalpy cycle (Born-Haber cycle)
6.5 Hesss law (SB p.155)

• Relate the various equations involved in a
  reaction

42
Importance of Hesss law
6.5 Hesss law (SB p.155)

• The enthalpy change of some chemical reactions
  cannot be determined directly because
• the reactions cannot be performed in the
  laboratory
• the reaction rates are too slow
• the reactions may involve the formation of side
  products

But the enthalpy change of such reactions can be
determined indirectly by applying Hesss Law.
43
Enthalpy change of formation of CO(g)
6.5 Hesss law (SB p.153)
-393.5 - (-283.0 )
-110.5 kJ mol-1
44
Enthalpy change of formation of CaCO3(s)
6.5 Hesss law (SB p.153)
45
Enthalpy change of hydration of MgSO4(s)
6.5 Hesss law (SB p.153)
aq
46
Calculations involving Standard Enthalpy Changes
of Reactions
47
Calculation of standard enthalpy change of
reaction from standard enthalpy changes of
formation
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.159)
48
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.159)
49
Calculation of standard enthalpy change of
formation from standard enthalpy changes of
combustion
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.162)
50
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.162)
51
6.7 Entropy change (SB p.164)
Entropy Change
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Entropy
6.7 Entropy change (SB p.164)

• A process is said to be spontaneous
• If no external forces are required to keep the
  process going
• The process may be physical change or a chemical
  change
• Example of spontaneous physical change cooling
  of hot water
• Example of spontaneous chemcial change burning
  of wood once the fire is started

53
6.7 Entropy change (SB p.164)
Entropy

• Exothermicity is the reason for the spontaneity
  of a process
• Some spontaneous changes are endothermic
• Examples Melting of ice, dissolution of ammonium
  nitrate in water

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6.7 Entropy change (SB p.164)
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6.7 Entropy change (SB p.165)
Entropy

• Entropy is a measure of the randomness or the
  degree of disorder of a system

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6.7 Entropy change (SB p.166)
Entropy change

• Entropy change means the change in the degree of
  disorder of a system
• ?S Sfinal - Sinitial

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6.7 Entropy change (SB p.166)
Positive entropy

• Increase in entropy
• Final state has a larger entropy that the initial
  state
• Example
• Ice (less entropy) ?? Water (more entropy)
• ?S Swater Sice ve

58
6.7 Entropy change (SB p.166)
Negative entropy

• Decrease in entropy
• Initial state has a larger entropy that the final
  state
• Example
• Water (more entropy) ?? Ice (less entropy)
• ?S Sice Swater -ve

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Free Energy Change
60
6.8 Free energy change (SB p.168)
Free energy change

• Entropy is temperature dependent
• At a higher temp, the entropy of the system is
  higher
• At a lower temp, the entropy of the system is
  lower

61
6.8 Free energy change (SB p.168)
Free energy

• Another driving force for a process
• Called free energy (G)
• G H TS
• where H is the enthalpy
• T is Kelvin temperature
• S is the entropy

62
6.8 Free energy change (SB p.168)
Free energy change

• ?G ?H T?S
• At a given temp, there are two driving forces for
  a process to occur
• Overall enthalpy of the system tends to be low
• Overall entropy of the system tends to be high

63
6.8 Free energy change (SB p.168)
Significance of the equation

• Process favoured by
• ?H -ve ?S ve
• Process not favoured by
• ?H ve ?S -ve

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6.8 Free energy change (SB p.168)
Significance of the equation

• A process is spontaneous or favourable when ?G is
  negative
• A process is not spontaneous or favourable as
  indicated when ?G is positive, but is spontaneous
  in the opposite direction

65
6.8 Free energy change (SB p.169)
How ?H and ?S affect the spontaneity of a process
?H ?S ?G ?H - T?S Result
-ve ve -ve Process is spontansous at all temepratures
ve -ve ve Process is not spontaneous at any temperature (reverse process is spontaneous at all temperatures)
66
6.8 Free energy change (SB p.170)
Effects of relative magnitudes of ?H and ?S on
the spontaneity of a process
?H ?S Condition ?G ?H - T?S Result
ve ve At high temp, T?S gt ?H ?G -ve Process is spontaneous at high temp
ve ve At low temp, ?H gt T?S ?G ve Process is not spontaneous
-ve -ve At high temp, T?S gt ?H ?G ve Process is not spontaneous
-ve -ve At low temp, ?H gt T?S ?G -ve Process is spontaneous at low temp
67
6.8 Free energy change (SB p.170)
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The END
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6.1 What is energetics? (SB p.140)
Back
Check Point 6-1
State whether the following processes are
exothermic or endothermic. (a) Melting of
ice. (b) Dissolution of table salt. (c)
Condensation of steam.
Answer
70
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.141)
Check Point 6-2

• State the difference between exothermic and
  endothermic reactions with respect to
• (i) the sign of ?H
• (ii) the heat change with the surroundings
• (iii) the total enthalpy of reactants and
  products.

Answer

• (i) Exothermic reactions ?H -ve endothermic
  reactions ?H ve
• (ii) Heat is given out to the surroundings in
  exothermic reactions whereas heat is taken in
  from the surroundings in endothermic reactions.
• (iii) In exothermic reactions, the total
  enthalpy of products is less than that of the
  reactants. In endothermic reactions, the total
  enthalpy is greater than that of the reactants.

71
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.141)
Check Point 6-2

• Draw an enthalpy level diagram for a reaction
  which is
• (i) endothermic, having a large activation
  energy.
• (ii) exothermic, having a small activation
  energy.

Answer
72
6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.141)
Check Point 6-2
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6.2 Enthalpy changes related to breaking and
forming of bonds (SB p.141)
Check Point 6-2
Back
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6.3 Standard enthalpy changes (SB p.147)
Check Point 6-3

1. Why must the condition burnt completely in
   oxygen be emphasized in the definition of
   standard enthalpy change of combustion?

Answer

1. If the substance is not completely burnt in
   excess oxygen, other products such as C(s) and
   CO(g) may be formed. The enthalpy change of
   combustion measured will not be accurate.

75
6.3 Standard enthalpy changes (SB p.147)
Check Point 6-3

• (b) The enthalpy change of the following reaction
  under standard conditions is 566.0 kJ.
• 2CO(g) O2(g) ?? 2CO2(g)
• What is the standard enthalpy change of
  combustion of carbon monoxide?

Answer
76
6.3 Standard enthalpy changes (SB p.147)
Check Point 6-3
(c) Enthalpy change of combustion of nitrogen or
enthalpy change of formation of nitrogen dioxide.
Answer
Back
77
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.149)
Example 6-4A
Determine the enthalpy change of neutralization
of 25 cm3 of 1.25 M hydrochloric acid and 25 cm3
of 1.25 M sodium hydroxide solution using the
following data Mass of calorimeter 100
g Initial temperature of acid 15.5 oC (288.5
K) Initial temperature of alkali 15.5 oC
(288.5 K) Final temperature of the reaction
mixture 21.6 oC (294.6 K) The specific heat
capacities of water and calorimeter are 4200 J
kg-1 K-1 and 800 J kg-1 K-1 respectively.
Answer
78
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.149)
Example 6-4A
Assume that the density of the reaction mixture
is the same as that of water, i.e. 1 g cm-3. Mass
of the reaction mixture (25 25) cm3 ? 1 g
cm-3 50 g 0.05 kg Heat given out (m1c1
m2c2) ?T (0.05 kg ?
4200 J kg-1 K-1 0.1 kg ? 800 J kg-1 K-1) ?
(294.6 288.5) K
1769 J H(aq) OH-(aq) ?? H2O(l) Number of
moles of HCl 1.25 mol dm-3 ? 25 ? 10-3 dm3
0.03125 mol Number of moles of NaOH 1.25 mol
dm-3 ? 25 ? 10-3 dm3 0.03125 mol Number of
moles of H2O formed 0.03125 mol
79
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.149)
Example 6-4A
Back
80
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.151)
Example 6-4B
Determine the enthalpy change of combustion of
ethanol using the following data Mass of spirit
lamp before experiment 45.24 g Mass of spirit
lamp after experiment 44.46 g Mass of water in
copper calorimeter 50 g Mass of copper
calorimeter without water 380 g Initial
temperature of water 18.5 oC (291.5 K) Final
temperature of water 39.4 oC (312.4 K) The
specific heat capacities of water and copper
calorimeter are 4200 J kg-1 K-1 and 2100 J kg-1
K-1 respectively.
Answer
81
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.151)
Example 6-4B
82
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.151)
Example 6-4B
Back
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6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.152)
Example 6-4C
0.02 mol of anhydrous ammonium chloride was added
to 45 g of water in a polystyrene cup to
determine the enthalpy change of solution of
anhydrous ammonium chloride. It is found that
there was a temperature drop from 24.5 oC to
23.0 oC in the solution. Given that the specific
heat capacity of water is 4200 J kg-1 K-1
and NH4Cl(s) aq ?? NH4Cl(aq) Calculate the
enthalpy change of solution of anhydrous ammonium
chloride. (Neglect the specific heat capacity of
the polystyrene cup.)
Answer
84
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.152)
Example 6-4C
Back
85
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4
(a) A student tried to determine the enthalpy
change of neutralization by putting 25.0 cm3 of
1.0 M HNO3 in a polystyrene cup and adding 25.0
cm3 of 1.0 M NH3 into it. The temperature rise
recorded was 3.11 oC. Given that the mass of the
polystyrene cup is 250 g, the specific heat
capacities of water and the polystyrene cup are
4200 J kg-1 K-1 and 800 J kg-1 K-1 respectively.
Determine the enthalpy change of neutralization
of nitric acid and aqueous ammonia. (Density of
water 1 g cm-3)
Answer
86
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4

• Heat evolved m1c1?T m2c2 ?T
• 0.050 kg ? 4200 J
  kg-1 K-1 ? 3.11 K 0.25 kg ?
  800 J kg-1 K-1 ? 3.11 K
• (653.1 622) J
• 1275.1 J
• No. of moles of HNO3 used 1.0 M ? 25 ? 10-3 m3
• 
  0.025 mol

87
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4
88
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4

• When 0.05 mol of silver nitrate was added to 50 g
  of water in a polystyrene cup, a temperature drop
  of 5.2 oC was recorded. Assuming that there was
  no heat absorption by the polystyrene cup,
  calculate the enthalpy change of solution of
  silver nitrate.
• (Specific heat capacity of water 4200 J kg-1
  K-1)

Answer
89
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4
90
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4
(c) A student used a calorimeter as shown in
Fig. 6-15 to determine the enthalpy change of
combustion of methanol. In the experiment, 1.60 g
of methanol was used and 50 g of water was
heated up, raising the temperature by 33.2 oC.
Given that the specific heat capacities of water
and copper calorimeter are 4200 J kg-1 K-1 and
2100 J kg-1 K-1 respectively and the mass of the
calorimeter is 400 g, calculate the enthalpy
change of combustion of methanol.
Answer
91
6.4 Experimental determination of enthalpy
changes by calorimetry (SB p.153)
Check Point 6-4
Back
92
6.5 Hesss law (SB p.158)
Check Point 6-5
Answer
93
6.5 Hesss law (SB p.158)
Check Point 6-5
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6.5 Hesss law (SB p.158)
Check Point 6-5
Answer
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6.5 Hesss law (SB p.158)
Check Point 6-5
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6.5 Hesss law (SB p.158)
Check Point 6-5
Answer
97
6.5 Hesss law (SB p.158)
Check Point 6-5
Back
98
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.159)
Example 6-6A
Answer
99
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.159)
Example 6-6A
Back
100
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.160)
Example 6-6B
Answer
101
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.160)
Example 6-6B
Back
102
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.160)
Example 6-6C
Answer
103
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.160)
Example 6-6C
Back
104
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.161)
Example 6-6D
Answer
105
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.161)
Example 6-6D
Back
106
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.162)
Example 6-6E
Answer
107
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.162)
Example 6-6E
108
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.163)
Back
Example 6-6E
109
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.163)
Example 6-6F
Answer
110
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.163)
Example 6-6F
111
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.163)
Back
Example 6-6F
112
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.164)
Check Point 6-6
Answer
113
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.164)
Check Point 6-6
114
6.6 Calculations involving standard enthalpy
changes of reactions (SB p.164)
Check Point 6-6
Answer
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6.6 Calculations involving standard enthalpy
changes of reactions (SB p.164)
Back
Check Point 6-6
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6.7 Entropy change (SB p.167)
Back
Check Point 6-7

• Predict whether the following changes or
  reactions involve an increase or a decrease in
  entropy.
• Dissolving salt in water to form salt solution
• Condensation of steam on a cold mirror
• Complete combustion of carbon
• Complete combustion of carbon monoxide
• Oxidation of sulphur dioxide to sulphur
  trioxide

Answer
117
6.8 Free energy change (SB p.170)
Back
Let's Think 1
In the process of changing of ice to water, at
what temperature do you think ?G equals 0?
Answer
?G equals 0 means that neither the forward nor
the reverse process is spontaneous. The system is
therefore in equilibrium. Melting point of ice is
0 oC (273 K) at which the process of changing ice
to water and the process of water turning to ice
are at equilibrium. At 0 oC, ?G of the processes
equals 0.
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6.8 Free energy change (SB p.170)
Check Point 6-8

• At what temperatures is the following process
  spontaneous at 1 atmosphere?
• Water ?? Steam
• What are the two driving forces that determine
  the spontaneity of a process?

Answer
119
6.8 Free energy change (SB p.170)
Check Point 6-8
Back

• State whether each of the following cases is
  spontaneous at all temperatures, not spontaneous
  at any temperature, spontaneous at high
  temperatures or spontaneous at low temperatures.
• (i) positive ?S and positive ?H
• (ii) positive ?S and negative ?H
• (iii) negative ?S and positive ?H
• (iv) negative ?S and negative ?H

1. Spontaneous at high temperatures
2. Spontaneous at all temperatures
3. Not spontaneous at any temperature
4. Spontaneous at low temperatures

Answer