Chemistry 8: Thermodynamics
A.
Enthalpy (H): Bond Energy (5.3 to 5.5, 8.8)
1. chemical reactions typically involve breaking bonds
between reactant atoms and forming new bonds
2. breaking bonds takes energy  chemical system gains
bond energy; surroundings lose energy (heat, etc.)
3. forming bonds releases energy  chemical system
loses energy, surroundings gain energy
4. change in energy called “change in enthalpy”—H
a. when energy required to break bonds > energy
released to form new bonds, +H (endothermic)
1. products at a higher energy state than
reactants (weaker bonds)
2. surroundings lose energy (cool down)
b. when energy required to break bonds < energy
released to form new bonds, –H (exothermic)
1. products at a lower energy state than
reactants (stronger bonds)
2. surroundings gain energy (heat up)
5. thermochemical equation
a. chemical equation with H
1. listed to the right of equation
2. included as reactant (endothermic) or product
(exothermic)
b. H can be used in dimensional analysis process
6. H from calorimetry
a. reactants are put in an insulated container filled
with water, where heat is exchanged between
reactants and water, but no heat is lost
b. by conservation of energy: Hreaction = –Qwater
1. Q = mcT for simple coffee cup calorimeter—
aqueous reactions
a. m = mass of solution ( water only)
b. c = specific heat of solution (usually same
as pure water—4.18 J/g•K)
c. T = change in temperature (Tf – Ti)
can stay in oC when T
2. Q = (C + mc)T for “bomb" calorimeter
a. C = “bomb constant” accounts for all
non-water components that change
temperature
b. all other letters are the same as the
simple calorimeter
7. H using bond energy (BE) data
Bond Energies in (kJ/mol)
Single
Multiple
H C N O
S
F Cl Br
I C=C 614
H 436 413 391 463 339 567 431 366 299 C=N 615
C
348 293 358 259 485 328 276 240 C=O 799
N
163 201
272 200 243
N=N 418
O
146
190 203
243 N=O 607
S
266 327 253 218
O=O 495
F
155 253 237
O=S 523
Cl
242 218 208 S=S 418
Br
193 175 CC 839
151 CN 891
I
CO 1072
NN 941
a. energy needed to break a bond (i.e. C–H) in a
diatomic, gaseous molecule, which contains the
bond type
1. is approximately the same for any molecule
2. affected by molecular bonding  only works
for gaseous species
3. positive value (+ BE) for breaking bonds
b. forming bonds (– BE)
c. H = BEbreaking – BEforming
Name __________________________
B.
Entropy (S): Disorder (19.2)
1. atoms/molecules have inherent disorder depending on
a. number of atoms—more internal motion = disorder
b. spacing of molecules—farther apart = disorder
c. speed of molecules—faster = disorder
2. predict increase in disorder for physical changes (+ S)
a. spread out: evaporation, diffusion and effusion
(solution: spread out solute and solvent (+S), but
bond solute-solvent (-S)  ?, but usually +S)
b. motion: melting, boiling
3. predict increase in disorder for chemical changes (+ S):
moles gaseous products > moles gaseous reactants
C. Thermodynamic Data (5.6 to 5.7, 19.4)
So (kJ/mol•K)
Species
Hfo (kJ/mol)
Al
0.0
+0.0283
Al3+
-531.0
-0.3217
Al2O3
-1675.7
+0.0509
1. standard heat of formation (Hfo) data
a. Ho for the formation of one mole of compound
from its elements at standard temperature (25oC)
Al: Al(s)  Al(s)  no reaction
Al3+: Al(s)  Al3+ + 3 eAl2O3: 2 Al(s) + 3/2 O2(g)  Al2O3(s)
b. Hfo for elements in natural state = 0.0 kJ/mol
c. more negative = more stable (harder to decompose)
2. standard entropy (So) data
a. amount of disorder compared to H+ (simplest form
of matter), which is zero by definition
b. listed in J/mol•K on AP exam, so you will have to
convert to kJ/mol•K for most calculations
3. calculations using the thermodynamic data chart
a. altering Hfo
1. opposite sign for the reverse reaction
C + 2 Cl2  CCl4 = –139.4 kJ
 CCl4  C + 2 Cl2 = +139.4 kJ
2. multiply by number of moles (coefficient)
1 mole CCl4= –139.4 kJ
 2 mole CCl4 = –278.8 kJ
b. calculate H for a reaction using Hfo
1. Hess’s Law: H for a multi-step reaction equals
the sum of H for each step
-(-74.8)
CH4(g)  C + 2 H2
-393.5
C + O2  CO2(g)
2(-241.8)
+ 2 H2 + O2  2 H2O(g)
CH4(g) + 2 O2  CO2(g) + 2 H2O(g) -802.3
2. H Ho = Hfoproducts – Hforeactants
c. calculate S for a reaction using So

S So = Soproducts – Soreactants
D. Gibbs Free Energy (G): Overall Energy State (19.5 to 19.6)
1. spontaneous chemical and physical changes occur
because the products are at a lower energy state
and/or more disordered
a. Gibbs free energy accounts for both H and S
b. change in free energy: G  Ho – TSo
2. determine if thermodynamically favorable (G < 0)
a. lower potential energy (-H)—chemical reactions
b. greater disorder (+S)—physical changes
c. threshold temperature (Tthreshold) when G = 0
Tthreshold = Ho/So
d. summary chart for determining when G < 0
- H
+ H
+ S
G < 0 for all T
G < 0 when T > Ho/So
-S G < 0 when T < Ho/So
G > 0 for all T
3. Go = –nFEo (in joules)
a. n: # e- in balanced redox equation
b. F: faraday = 96,500 C/mol e-
d.
Temperature (oC)
Heat of Reaction Lab—Use calorimetry to determine H
for a series of reactions, compare the results with
thermodynamic data, and combine the results to verify
Hess' law.
Part 1 (Determine C) Heat about 75 mL of water to about
70oC. Place a Styrofoam cup in a 250-mL beaker. Add 50.0
mL cold tap water to the cup. Record the temperature TC.
Measure out 50.0 mL of the hot water and place in a second
Styrofoam cup. Record the temperature TH. Pour the hot
water into the cold water, cover the cup, insert the
thermometer in the hole, and mix gently. Record the
temperature every 20 seconds for 3 minutes. Discard.
a. (1) Record the temperatures.
TC
TH
time (s) 20
40 60 80 100 120 140 160 180
To C
(2) Graph the temperature vs. time data. Draw a best
fit straight line (use a ruler).
20
60
100
140
180
Time (s)
(3) Use the y-intercept to determine Tmix.
Tmix (y-intercept)
b. Calculate the following.
(1) Average of the hot and cold temperatures.
Tav = (TH + TC)/2
(2) Heat lost from the water.
QL = mc(Tav – Tmix)
(3) C.
C = QL/(Tmix – TC)
Part 2 (OH- + H+  H2O) Place a Styrofoam cup in a 250
mL beaker. Add 50.0 mL of 3.00 M NaOH. Record the
temperature To. Pour 50.0 mL of 3.00 M HCl into the NaOH,
cover, insert the thermometer, and mix gently. Record the
temperature every 20 seconds for 3 minutes. Discard.
c. (1) Record the temperatures.
To
time (s) 20
40 60 80 100 120 140 160 180
ToC
(2) Graph the temperature vs. time data. Draw a best
fit straight line (use a ruler).
T (K)
Qwater (kJ)
Hreaction/mole
e.
Calculate Hreaction per mole of reactant based on Hfo.
OH-(aq) + H+(aq)  H2O(l)
H
%
Part 3 (NH4+ + OH-  NH3(aq) + H2O) Place a Styrofoam
cup in a 250 mL beaker. Add 50.0 mL of 3.00 M NH4Cl.
Record the temperature To. Pour 50.0 mL of 3.00 M NaOH
into the NH4Cl, cover, insert the thermometer, and mix
gently. Record the temperature every 20 seconds for 3
minutes. Discard.
f. (1) Record the temperatures.
To
time (s) 20
40 60 80 100 120 140 160 180
ToC
(2) Graph the temperature vs. time data. Draw a best
fit straight line (use a ruler).
20
100
140
Time (s)
(3) Use the y-intercept to determine Tmix.
Tmix (y-intercept)
g.
60
180
Calculate Hreaction per mole of reactant based on the
calorimetry data.
T (K)
Qwater (kJ)
Hreaction/mole
h.
Calculate Hreaction per mole of reactant based on Hfo.
NH4+(aq) + OH-(aq)  NH3(aq) + H2O(l)
H
%
Temperature (oC)
1.
Calculate Hreaction per mole of reactant based on the
calorimetry data.
Temperature (oC)
Experiments
20
60
100
140
Time (s)
(3) Use the y-intercept to determine Tmix.
Tmix (y-intercept)
180
Part 4 (NH3(aq) + H+  NH4+) Place a Styrofoam cup in a
250 mL beaker. Add 50.0 mL of 3.00 M NH3. Record the
temperature To. Pour 50.0 mL of 3.00 M HCl into the NH3,
cover, insert the thermometer, and mix gently. Record the
temperature every 20 seconds for 3 minutes. Discard.
i. (1) Record the temperatures.
To
time (s) 20
40 60 80 100 120 140 160 180
ToC
4.
Temperature (oC)
(2) Graph the temperature vs. time data. Draw a best
fit straight line (use a ruler).
20
100
140
Time (s)
(3) Use the y-intercept to determine Tmix.
Tmix (y-intercept)
j.
60
b.
the number of moles of MgSO4 dissolved.
c.
H (in kJ) for the dissolving of one mole of MgSO4.
180
Calculate Hreaction per mole of reactant based on the
calorimetry data.
5.
H2(g) + F2(g)  2 HF(g)
Estimate H for the reaction using the bond energy values.
6.
C2H2(g) + 2 H2(g)  C2H6(g)
Estimate H for the reaction using the bond energy values.
7.
A bomb calorimeter with a constant of 921 J/oC contains
1,000 g of water. The combustion of 1.00 g of ethene
(C2H4) increases the temperature 9.3oC. Determine
a. Qwater for the combustion process.
T (K)
Qwater (kJ)
Hreaction/mole
k.
12.8 g of MgSO4 is dissolved in 250. g of H2O in a coffee
cup calorimeter. The temperature of the solution increases
from 23.8oC to 33.1oC. Determine
a. Qwater for the solution process.
Calculate Hreaction per mole of reactant based on Hfo.
NH3(aq) + H+(aq)  NH4+(aq)
H
b.
the number of moles of ethene reacted.
c.
H (in kJ) for the combustion of one mole of C2H4.
d.
Write the equation for the combustion of ethene (C2H4).
e.
Calculate H using bond energies.
%
l.
Show that the chemical equations and H from
Part 2 = Part 3 + Part 4.
Practice Problems
1.
a.
b.
2.
8.
How many grams of iron are needed to generate
1.00 x 104 kJ of heat?
CaSO4(s) + CO2(g)  CaCO3(s) + SO3(g)
H = 224 kJ
a. How much heat is absorbed when 10.0 g CaSO4 react.
b.
3.
A. Enthalpy
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
H = -1640 kJ
How much heat is released to produce 10.0 g Fe2O3?
9.
B. Entropy
Predict whether S > 0, S < 0 or S 0.
>0
 0 < 0
Melting ice at 0oC
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
Distilling alcohol-water mixture
C. Thermodynamic Data
2 Na2O2(s) + 4 HCl(g)  4 NaCl(s) + 2 H2O(l) + O2(g)
Determine H from the thermochemical reactions below.
2 Na2O2(s) + 2 H2O(l)  4 NaOH(s) + O2(g) H1 = -126 kJ
NaOH(s) + HCl(g) NaCl(s) + H2O(l)
H2 = -179 kJ
How much CaCO3 is produced when 500. kJ is absorbed?
10.
C2H2(g) + 5 N2O(g)  2 CO2(g) + H2O(g) + 5 N2(g)
Determine H from the thermochemical equations below.
2 C2H2(g) + 5 O2(g) 4 CO2(g) + 2 H2O(g) H1 = -2512 kJ
N2(g) + ½ O2(g)  N2O(g)
H2 = 104 kJ
11.
NO(g) + O(g)  NO2(g)
Determine H for the above reaction using the following
thermochemical equations.
NO(g) + O3(g)  NO2(g) + O2(g)
H1 = -198.9 kJ
O3(g)  3/2 O2(g)
H2 = -142.3 kJ
O2(g)  2 O(g)
H3 = 495.0 kJ
When 1.51 g of NH4Cl are dissolved in 100. g of water the
temperature drops 1.00oC. Determine
a. Qwater for the solution process.
b.
the number of moles of NH4Cl dissolved.
c.
H (in kJ) for the dissolving of one mole of NH4Cl.
12. a.
b.
Write the equation for the combustion of methanol,
CH3OH(l). (other reactants and products are gaseous).
b.
Calculate So.
Calculate H using Hfo values.
c.
Calculate G at 20oC.
d.
What is the spontaneous temperature range?
1.00 g of methanol is burned in a bomb calorimeter that
contains 1200 g of water. The temperature increases 3.4 K.
c. Calculate the heat generated by the combustion reaction.
d.
13.
Calculate the calorimeter constant of the bomb.
Ca(s) + SO3(g) + 2 H2O(l)  CaSO3•2 H2O(s)
H = -795 kJ and S = -0.2535 kJ/K for the reaction.
a. Calculate Hfo for CaSO3•2 H2O.
b.
18. When H2SO4(l) is dissolved in water, the temperature of
the mixture increases. Predict the sign of H, S and G
for this process (justify your answer).
+/–
Justification
H
S
G
19.
C2H5OH(l)  C2H5OH(g)
Calculate the boiling point (threshold temperature) given
the information: H = 37.95 kJ and S = 0.1078 kJ/K.
Calculate So for CaSO3•2 H2O.
Practice Multiple Choice
D. Gibbs Free Energy
14. Cu(s) + 4 H+(aq) + 2 NO3-(aq)  Cu2+(aq) + 2 NO2(g) + 2 H2O(l).
a. Calculate Ho using Hfo values.
15.
b.
Calculate So using So values.
c.
Calculate Go at 25oC.
Briefly explain why the answer is correct in the space provided.
1.
I2(g) + 3 Cl2(g)  2 ICl3(g)
What is Ho, in kJ, for the reaction represented above?
Bond
CI–CI
I–I
I–Cl
Bond Energy (kJ/mole)
150
240
210
(A) - 870
(B) - 390
(C) +180
(D) + 450
2.
C2H4(g) + 3 O2(g)  2 CO2(g) + 2 H2O(g)
For the reaction, H is -1,300 kJ. What is the value of H,
in kJ, if the combustion produced liquid water rather than
water vapor? (H for H2O(l)  H2O(g) is 45 kJ/mol)
(A) -1,300 (B) -1,210 (C) -1,345 (D) -1,390
3.
CH4 (g) + 2 O2(g)  CO2(g) + 2 H2O(l) Ho = -900 kJ
What is the standard heat of formation of CH4, in kJ/mol,
as calculated from the data below?
(HfoH2O = -300 kJ/mol, HfoCO2 = -400 kJ/mol)
(A) -200
(B) -100
(C) 100
(D) 200
4.
H2(g) + ½ O2(g)  H2O(l)
Ho = x
2 Na(s) + ½ O2(g)  Na2O(s)
Ho = y
Na(s) + ½ O2(g) + ½ H2(g)  NaOH(s)
Ho = z
What is H for the reaction: Na2O(s) + H2O(l)  2 NaOH(s)
(A) x + y + z (B) x + y – z (C) x + y - 2z
(D) 2z - x - y
5.
Which is true when ice melts at its normal melting point?
(A) H < 0, S > 0, G = 0 (B) H < 0, S < 0, G > 0
(C) H > 0, S < 0, G < 0 (D) H > 0, S > 0, G = 0
6.
Which of the following reactions has the largest positive
value of S per mole of Cl2?
(A) H2(g) + Cl2(g)  2 HCl(g)
(B) Cl2(g) + O2(g)  Cl2O(g)
(C) Mg(s) + Cl2(g)  MgCl2(s)
(D) 2 NH4Cl(s)  4 H2(g) + Cl2(g)
NH4NO3(s)  NH4+(aq) + NO3-(aq)
Determine the following for the above reaction.
a. Is the reaction exothermic or endothermic?
b.
Is there an increase or decrease in entropy?
c.
Is the reaction spontaneous at 25oC?
a.
2 SO2(g) + O2(g)  2 SO3(g)
Calculate Ho.
b.
Calculate So.
16.
c.
Calculate G at 400 K.
d.
Determine the temperature range where the reaction
is spontaneous.
17.
a.
C2H5OH(l) + 3 O2(g)  2 CO2(g) + 3 H2O(l)
Calculate Ho.
7.
Ice is added to hot water in an insulated container, which is
then sealed. What has happened to the total energy and
the total entropy when the system reaches equilibrium?
(A) Energy and entropy remain constant
(B) Energy remains constant, entropy decreases
(C) Energy remains constant, entropy increases
(D) Energy decreases, entropy increases
c.

d.
What is the spontaneous temperature range?

2.
N2(g) + 3 H2(g)  2 NH3(g)
The above reaction is thermodynamically spontaneous at
298 K, but becomes nonspontaneous at higher
temperatures. Which of the following is true at 298 K?
(A) G, H, and S are all positive.
(B) G, H, and S are all negative.
(C) G and H are negative, but S is positive.
(D) G and S are negative, but H is positive.
8.
Calculate Go for the combustion reaction at 25oC.
Consider the synthesis reaction: N2(g) + 3 F2(g)  2 NF3(g)
(Ho298 = -264 kJ, So298 = -278 J K-1)
o
a. Calculate G 298 for the reaction.

b.
For what temperature range is the reaction spontaneous?

c.
Calculate the heat released when 0.256 mol of NF3(g)
is formed from N2(g) and F2(g) at 1.00 atm and 298 K.

3 C2H2(g)  C6H6(g)
What is the standard enthalpy change, Ho, for the
reaction represented above?
(HfoC2H2 is 230 kJ•mol-1; HfoC6H6 is 80 kJ•mol-1)
(A) -610 kJ (B) 150 kJ (C) -770 kJ (D) 610 kJ
9.
10. When solutions of NH4SCN and Ba(OH)2 are mixed in a
closed container, the temperature drops and a gas is
produced. Which of the following indicates the correct
signs for G, H, and S for the process?
(A) –G –H –S (B) –G +H –S
(C) –G +H +S (D) +G –H +S
11.
X(s)  X(l)
Which of the following is true for any substance undergoing
the process represented above at its normal melting point?
(A) S < 0
(B) H = 0
(C) H = TG
(D) H = TS
12. For a reaction, Ho = -150 kg/mol and So = -50 J/mol•K.
Which statement is true about this reaction?
(A) It is spontaneous at high temperature only.
(B) It is spontaneous at low temperature only.
(C) It is spontaneous at all temperatures.
(D) It is non-spontaneous at all temperatures.
Practice Free Response
1.
Consider the combustion of butanoic acid at 25oC:
HC4H7CO2(l) + 5 O2(g)  4 CO2(g) + 4 H2O(l) Ho= -2,183.5 kJ
So (kJ/mol•K)
Substance
Hfo (kJ/mol)
CO2(g)
-393.5
0.2136
H2O(l)
-285.8
0.0699
O2(g)
0.0
0.2050
C3H7COOH(l)
?
0.2263
a. Calculate Hfo, for butanoic acid.
b.

Calculate So for the combustion reaction at 25oC.
d.
Calculate the F–F bond energy using the information
above and the bond energies
(NN = 946 kJ/mol, N–F = 272 kJ/mol).

3.
The dissolving of AgNO3(s) in water is represented by the
equation: AgNO3(s)  Ag+(aq) + NO3-(aq)
a. Is G positive, negative, or zero? Justify your answer.

b.
The solution cools when AgNO3(s) is dissolved. Is H
positive, negative or zero? Justify your answer.

c.
Is S positive, negative, or zero? Justify your answer.

4.
Consider the thermochemical equation:
C2H6(g) + 7/2 O2(g)  2 CO2(g) + 3 H2O(l) H = -1559.7 kJ
a. Calculate H for the thermochemical equations:
2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(l)
2 CO2(g) + 3 H2O(l) C2H6(g) + 7/2 O2(g) 
b. The heat of vaporization of H2O(l) is +44.0 kJ/mol.
Calculate H for the equation:
C2H6(g) + 7/2 O2(g)  2 CO2(g) + 3 H2O(g)
c.
The heat of formation of CO2(g) and H2O(l) are -393.5
kJ/mol and -285.8 kJ/mol. Calculate Hfo of C2H6(g).
d.
How much heat is evolved when 1.00 g of C2H6(g) is
burned to give CO2(g) + H2O(l) in an open container?
e.
What is the bomb constant C if the change in
temperature is 13.13oC when 1.00 g of C2H6 reacts in
the bomb calorimeter that contains 250 g H2O?