Download Chemical Thermodynamics (with Thermochemistry) Addresses the

ENERGY (THERMOCHEMISTRY)______ Ch 1.5, 6, 9.10, 11.5-11.7, 12.3
Thermochemistry ̶ Prediction and measurement of energy transfer, in the form of heat, that accompanies
chemical and physical processes.
Chemical Thermodynamics (with Thermochemistry) Addresses the Questions
1. Will a particular reaction occur?
2. If it does occur then
a) what energy changes and transfers are involved?
b) to what extent?
Energy (capacity to do work and/or cause heat transfer)
kinetic (motion)
potential (position, chemical composition)
energy can be transferred from one form to another
interconversion of potential and kinetic energy
interconversion of work and heat
heat, q (E transferred between two bodies at different T)
work, w (E used to cause motion against an external force)
1st Law of Thermodynamics ̶ Conservation of Energy: Energy can neither be
created nor destroyed, only converted from one form to another.
internal energy, E
ΔE = q + w
̶ heat adsorbed by system and work done on system
FIG I ̶ Energy Transfer Between a System and its Surroundings
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Definitions
universe
system
surroundings
system – characterized with respect to transfers of
matter/energy with surroundings
open
closed
isolated
surroundings
boundaries (walls) – characterized with respect to energy
transfers in the form of heat
adiabatic
diathermal
rigid/nonrigid
WHY? ………………………………………. ability to predict
REACTANTS
initial state
PRODUCTS
final state
prediction => state function => energy (ΔE?) => heat and
work
state function
A State Function: Change in altitude depends only on the
difference between initial and final values, not on the path.
GOAL: if we can find the correct state function we
can use it to predict the progress of a chemical reaction.
Is Work the Desired State Function? (consider work in an ideal gas expansion)
FIG II - Gas Confined to a Cylinder
with a Movable Piston
FIG III - Calculating the Work
I
)
I
f
4
T
=
O
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EX 1. For the combustion of ethanol at room temperature in the open air
CH3CH2OH(l) + O2(g) → CO2(g) + H2O(l)
What is the work done by the reaction? What is doing the work? What assumptions are you making?
Is Heat the Desired State Function? (consider a beaker of hot or cold water)
FIG IV - Heat Transfer Between a Beaker of Water and its Surroundings
Working with Heat Transfer- Calorimetry: Measurement of quantities of heat
heat at constant pressure, qP
heat at constant volume, qV
exothermic, endothermic
adiabatic (Greek, adiabatos – impassable)
thermal equilibrium
heat capacity (rule of Dulong and Petit)
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Heat Capacities of Some Substances (Table 6.4): cP = molar mass × cs
Substance
Specific Heat
cs (J K-1 g-1)
Hg(l)
0.1393
Al(s)
Fe(s)
Cu(s)
0.900
0.4493
0.3849
He(g)
H(g)
H2(g)
5.193
20.62
14.30
Molar Mass
(g mol-1)
200.59
26.982
55.845
63.546
4.0026
1.0079
2.0158
N2(g)
CH4(g)
CaCO3(s)
CH3COOH(l)
1.04
2.22
28.014
16.043
0.818
2.05
100.086
60.052
H2O(s)
H2O(l)
H2O(g)
2.1
4.184
2.0
18.015
18.015
18.015
Molar Heat Capacity
cP ( J K-1 mol-1)
27.94
24.3
25.09
24.46
20.79
20.78
28.82
29.14
35.6
81.9
123.1
38
75.38
36
EX 2. 6.42 g of He absorbs 500.0 kJ of heat (at constant pressure) and the temperature increases by
15.0 K.
What is CP?
What is cs?
What is cP?
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Calculations with Heat Transfer
calorimeter
EX 3. 5.00 g of Fe (cs = 0.449 J K-1 g-1) at 400 K is put into 10.0 g of water (cs = 4.18 J K-1g-1) at 290 K.
a) Ignoring the heat capacity of the calorimeter, what is the final temperature?
b)
If the calorimeter has a heat capacity of 500 J K-1 what is the final temperature?
EX 4. A styrofoam cup calorimeter contains 150 g of water. Generation of 1430 J of heat inside the calori­
meter causes the temperature to increase by 1.93°C. What is Ccal?
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Enthalpy (Greek, enthalpein – to warm in), H (heat at constant pressure, qp)
enthalpy of reaction
N2 (g) + 3 H2 (g) → 2 NH3(g)
per
2 N2 (g) + 6 H2 (g) → 4 NH3 (g)
2 NH3
ΔH = -91. 8 kJ
one mole of N2
three moles of H2
two moles of NH3
→ N2 (g) + 3 H2(g)
ΔH = -2(91.8) kJ
ΔH = +91.8 kJ
enthalpy of phase changes
FIG V- Energy Changes Accompanying Phase Changes
(Greek, phasis- appearance)
Common phase transitions involve a transfer
of heat between a system and its surroundings
at constant temperature and pressure.
11.5 vaporization g → l (condensation l → g)
11.6 fusion (melting) s → l (freezing l → s)
11.7 sublimation s → g (deposition g → s)
I
Internal Energy, E (heat at constant volume, qV)
Relation between ΔH and ΔE
EX 5. At atmospheric pressure and 25°C the detonation of nitroglycerin
4 C 3 H 5 ( N O 3 ) 3 ( l)
→
6 N 2 ( g ) + 1 0 H 2 O( g ) + 1 2 C O 2 ( g ) + O 2 ( g )
releases 5720 kJ of heat. What is the internal energy change for this reaction?
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Experimental Calorimetry
heat at constant pressure, qP
FIG VI - Coffee Cup Calorimeter
heat at constant volume, qV
FIG VII - An Adiabatic Bomb Calorimeter
EX 6. 0.1584 g of benzoic acid (C6H5COOH) are combusted in a constant volume bomb calorimeter. The
temperature of the calorimeter rises by 2.54°C. ΔE for the combustion of benzoic acid is -26.38 kJ g-1.
Determine the heat capacity of the calorimeter.
Hess's Law - Extremely USEFUL
combining enthalpies of reaction –
enthalpy is a state function
Hess’s Law: The
change in enthalpy
for a stepwise process is the sum of
the changes in enthalpy of the steps.
A + 2B → C
C → 2D
application to chemical reactions
EX 7. Find the enthalpy change that accompanies the combustion of graphite to carbon monoxide from:
ΔH = -393.5 kJ mol-1
1) C(s, gr) + O2(g) → CO 2 (g)
ΔH = 283.0 kJ mol-1
2) CO 2 (g) → CO(g) + 1/2 O 2 (g)
C(s, gr) + 1/2 O 2 (g)
→
CO(g)
ΔH = ?
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FIG VIII- Enthalpy Diagrams for Chemical Reactions
application to phase changes
FIG IX -Enthalpy Change upon Cooling 2.5 mol of Water from 130°C to -40°C
EX 8. How much heat is required to lower the temperature of 2.5 mol of water from 130°C to -40°C? For water:
cP(s), cP(l), cP(g) = 37.6, 75.38, 33.1 J mol-1 K-1 , ΔHvap, ΔHfus = 40.7, 6.02 kJ mol-l
overall: H2O(g, 130°C) → H2O(s, -40°C):
a)
H2O(g, 130°C) → H2O(g, 100°C) :
b)
H2O(g, 100°C) → H2O(l, 100°C) :
c)
H2O(l, 100°C) → H2O(l, °C) :
d)
H2O(l, °C) → H2O(s, -0°C) :
e)
H2O(s, 0°C) → H2O(s, -40°C) :
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"Standard" Enthalpies
standard states – stable form (allotrope) at P = 1 atm and specified temperature (usually T = 25°C)
pure solid
pure liquid
gas – ideal gas behavior
1 molar (1 M) solution – ideal solution behavior
standard enthalpy of formation, ∆Hfo, of a compound:
enthalpy change for the reaction forming one
mole of the compound from its elements in their standard states. ∆Hfo (element in standard state) = 0
EX 9. Find the enthalpy change that accompanies the combustion of methane to carbon dioxide from enthalpies of
formation (∆Hfo in Appendix IIB)
defining formation reaction
∆Hfo (kJ mol-1)
-74.81
1) C(s, gr) + 2 H2(g) → CH4(g)
O2(g)
→
O2(g)
3) C(s, gr) + O2(g)
→
CO2(g)
-393.51
4) H2(g) + 1/2 O2(g)
→
H2O (l)
-285.83
2)
0
TARGET: CH4(g) + 2 O2(g)
→
CO2(g) + 2 H2O(l)
long way:
short way:
FOR ANY REACTION:
∆Horx = ∑.nprod ∆Hof (products) – ∑.nreact ∆Hof (reactants)
=> products - reactants
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application to a) chemical reactions and b) phase changes
EX 10. For the combustion of ethanol (unbalanced)
CH 3 CH 2 O H( l) + O 2 ( g )
→
CO 2 ( g ) + H 2 O( l )
a) Determine the standard enthalpy of the reaction given that ∆Hfo for CH 3CH 2OH(l), CO 2(g), and H 2O(l) are
-277.0, -393.509, and -285.83, respectively.
b) What is the standard enthalpy change when all reactants and products are gases? At 25°C, ΔHvap(H 2O)
= 44.01 kJ mol-1 and ΔHvap (CH 3CH 2OH) = 42.59 kJ mol-1.
application to combustion reaction/ignoring the work term
EX 11. Determine ∆Hof of benzoic acid (C 6 H 5 C OO H). ΔE o for the combustion of benzoic acid is -26.38 kJ
g-1,∆Hof [CO 2(g)] = -393.5 kJ mol-1 ∆Hof [H 2O(1)] = -285.840 kJ mol-1.
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Average Bond Enthalpies (Chapter 9.10- Bond Energies)
Heat, at standard conditions, absorbed by a system to break chemical bonds in the gas phase.
Bond Energy of a Homonuclear Diatomic
bond enthalpy of chlorine gas (measured, 243 kJ/mol):
Cl2(g)
½ Cl2(g)
→ 2 Cl(g)
→ Cl(g)
ΔHo = ?
ΔHo = 121.68
2 × 121.68 = 243.46, close to 243
Bond Energy of Methane
1) atomization enthalpy of CH4 (energy to break all bonds and form free atoms, atomize, in gas phase):
CH4(g)
→
ΔHo = 1663 kJ/mol
per CH, 1663/4 = 415
C(g) + 4 H(g)
2) enthalpy of a C – H bond in CH4 (experimental results):
CH4(g)
CH3(g)
CH2(g)
CH(g)
→ CH3(g) + H(g)
→ CH2(g) + H(g)
→ CH(g) + H(g)
→ C(g) + H(g)
ΔHo = 439 kJ/mol
ΔHo = 453
ΔHo = 425
ΔHo = 339
average = 1656/4 = 414
3) enthalpy of a C – H bond:
CH4(g)
C2H6(g)
CHF3(g)
CHBr3(g)
→
→
→
→
ΔHo = 439 kJ/mol
ΔHo = 410
ΔHo = 429
ΔHo = 377
above 4 within 8% of average
CH3(g) + H(g)
C2H5(g) + H(g)
CF3(g) + H(g)
CBr3(g) + H(g)
Table 9-3. Average Bond Enthalpies
Single Bonds
H
C
N
O
F
Cl
Br
I
S
Si
H
436
412
388
463
565
431
366
299
338
318
C
348
305
360
484
338
276
238
259
N
O
F
163
157
270
200
146
185
203
158
254
496
Cl
Br
243
219
210
250
193
178
212
I
Si
S
151
264
226
466
Multiple Bonds
C=C
C≡C
O2
612
812
497
N=N
N≡N
CO 2
409
946
799
C=N
C≡N
613
890
C=O
C≡O
743
1046
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FIG X. Estimate ΔHo for:
2)
1)
-676
4)
-824
5)
-1130
6)
872
968
PRODUCTS
REACTANTS
3)
CC12F2 + 2 H 2 → CH 2Cl 2 + 2 HF
676
2516
-2630
=> reactants - products
EX 12. Use average bond enthalpies to estimate ΔHo for
2 CH 2 = CH – CH 3 (g) + 2 NH 3 (g) + 3 O 2 (g)
→
2 CH 2 = CH – CN(g) + 6 H 2O(g)