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HONORS CHEMISTRY
HARVARD-WESTLAKE
UNIT 5
Thermodynamics
What do you do with a dead chemist?
Barium
UNIT 5
1
Reference Sheet
Specific heat of water: 4.184 J/g°C
IA
VIIIA
H
2.1
Electronegativities of the Elements
IIA
IIIA
IVA
VA
VIA
VIIA
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
K
IIIB
IVB
VB
VIB
VIIB
VIII
IB
IIB
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
1.5
1.8
2.1
2.5
3.0
Ni
Cu
Zn
Ga
Ge
As
Se
Br
0.8
1.0
1.3
1.5
1.6
1.6
1.5
1.8
1.9
1.9
1.9
1.6
1.6
1.8
2.0
2.4
2.8
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
0.8
1.0
1.6
1.8
Ba
1.2
La-Lu
1.4
Cs
Hf
Ta
W
1.9
2.2
2.2
2.2
1.9
1.7
1.7
1.8
1.9
2.1
2.5
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
0.7
0.9
1.01.2
At
1.3
1.5
1.7
1.9
2.2
2.2
2.2
2.4
1.9
1.8
1.9
1.9
2.0
2.2
Fr
Ra
0.7
0.9
Table of Bond Energies
Bond
*H–H
H–N
H–O
H–S
H–P
*H–F
*H–Cl
*H–Br
*H–I
C–H
C–C
C=C
C≡C
Bond Energy
(kJ/mol)
436.4
393
460
368
326
568.2
431.9
366.1
298.3
414
347
620
812
Bond
C–Cl
C–N
C=N
C≡N
C–O
C=O
C=O (CO2)
C–P
C–S
C=S
N–N
N=N
*N≡N
Bond Energy
(kJ/mol)
339
276
615
891
351
745
799
263
255
477
193
418
941.4
Bond
N–O
N–P
O–O
*O=O
O=S
P–P
P=P
S–S
S=S
*F–F
*Cl–Cl
*Br–Br
*I–I
Bond Energy
(kJ/mol)
176
209
142
498.7
469
197
489
268
352
156.9
242.7
192.5
151.0
*Bolded values represent exact values for bonds which occur in diatomic molecules.
All other values represent averages of all instances of those bonds in larger molecules.
UNIT 5
2
Table of Thermodynamic Values: ∆H˚f, ∆G˚f, and S˚ at 1 atm and 25 ˚C
∆H˚f (kJ/mol)
∆G˚f (kJ/mol)
S˚ (J/K•mol)
Al (s)
Al3+ (aq)
Al2O3 (s)
0
-524.7
-1669.8
0
-481.2
-1576.4
28.3
-313.38
50.99
Br2 (l)
Br- (aq)
HBr (g)
0
-120.9
-36.2
0
-102.8
-53.2
152.3
80.7
198.48
C (graphite)
C (diamond)
CO (g)
CO2 (g)
CO2 (aq)
CO32- (aq)
HCO3- (aq)
H2CO3 (aq)
CS2 (g)
CS2 (l)
HCN (aq)
CN- (aq)
(NH2)2CO (s)
(NH2)2CO (aq)
0
1.90
-110.5
-393.5
-412.9
-676.3
-691.1
-399.7
115.3
87.3
105.4
151.0
-333.19
-319.2
0
2.87
-137.3
-394.4
-386.2
-528.1
-587.1
-623.2
65.1
63.6
112.1
165.69
-197.15
-203.84
5.69
2.4
197.9
213.6
121.3
-53.1
94.98
187.4
237.8
151.0
128.9
117.99
104.6
173.85
Ca (s)
Ca2+ (aq)
CaO (s)
Ca(OH)2 (s)
CaF2 (s)
CaCl2 (s)
CaSO4 (s)
CaCO3 (s)
0
-542.96
-635.6
-986.6
-1214.6
-794.96
-1432.69
-1206.9
0
-553.0
-604.2
-896.8
-1161.9
-750.19
-1320.3
-1128.8
41.6
-55.2
39.8
76.2
68.87
113.8
106.69
92.9
Cl2 (g)
Cl- (aq)
HCl (g)
0
-167.2
-92.3
0
-131.2
-95.27
223.0
56.5
187.0
Cu (s)
Cu+ (aq)
Cu2+ (aq)
CuO (s)
Cu2O (s)
CuCl
0
51.88
64.39
-155.2
-166.69
-134.7
0
50.2
64.98
-127.2
-146.36
-118.8
33.3
-26.4
-99.6
43.5
100.8
91.6
Substance
UNIT 5
3
Substance
∆H˚f (kJ/mol)
∆G˚f (kJ/mol)
S˚ (J/K•mol)
CuCl2 (s)
CuS (s)
CuSO4 (s)
-205.85
-48.5
-769.86
?
-49.0
-661.9
?
66.5
113.39
F2 (g)
F- (aq)
HF (g)
0
-329.1
-271.6
0
-276.48
-270.7
203.34
-9.6
173.5
Fe (s)
Fe2+ (aq)
Fe3+ (aq)
Fe2O3 (s)
Fe(OH)2 (s)
Fe(OH)3 (s)
0
-87.86
-47.7
-822.2
-568.19
-824.25
0
-84.9
-10.5
-741.0
-483.55
?
27.2
-113.39
-293.3
90.0
79.5
?
H (g)
H2 (g)
H+ (aq)
OH- (aq)
H2O (g)
H2O (l)
H2O2 (l)
218.2
0
0
-229.94
-241.8
-285.8
-187.6
203.2
0
0
-157.30
-228.6
-237.2
-118.1
114.6
131.0
0
-10.5
188.7
69.9
?
0
-55.9
25.9
0
51.67
1.30
116.7
109.37
206.3
K (s)
K+ (aq)
KOH (s)
KCl (s)
KClO3 (s)
KClO4 (s)
KBr (s)
KI (s)
KNO3 (s)
0
-251.2
-425.85
-435.87
-391.20
-433.46
-392.17
-327.65
-492.7
0
-282.28
?
-408.3
-289.9
-304.18
-379.2
-322.29
-393.1
63.6
102.5
?
82.68
142.97
151.0
96.4
104.35
132.9
Li (s)
Li+ (aq)
Li2O (s)
LiOH (s)
0
-278.46
-595.8
-487.2
0
-293.8
?
-443.9
28.0
14.2
?
50.2
Mg (s)
Mg2+ (aq)
0
-461.96
0
-456.0
32.5
-117.99
I2 (s)
I- (aq)
HI (g)
UNIT 5
4
∆H˚f (kJ/mol)
∆G˚f (kJ/mol)
S˚ (J/K•mol)
MgO (s)
Mg(OH)2 (s)
MgCl2 (s)
MgSO4 (s)
MgCO3 (s)
-601.8
-924.66
-641.8
-1278.2
-1112.9
-569.6
-833.75
-592.3
-1173.6
-1029.3
26.78
63.1
89.5
91.6
65.69
N2 (g)
N3- (aq)
NH3 (g)
NH4+ (aq)
NH4Cl (s)
NH3 (aq)
N2H4 (l)
NO (g)
NO2 (g)
N2O4 (g)
N2O (g)
HNO3 (aq)
0
245.18
-46.3
-132.80
-315.39
-80.3
50.4
90.4
33.85
9.66
81.56
-207.4
0
?
-16.6
-79.5
-203.89
-263.76
?
86.7
51.8
98.29
103.6
-111.3
191.5
?
193.0
112.8
94.56
111.3
?
210.6
240.46
304.3
219.99
146.4
Na (s)
Na+ (aq)
Na2O (s)
NaCl (s)
NaI (s)
Na2SO4 (s)
NaNO3 (s)
Na2CO3 (s)
NaHCO3 (s)
0
-239.66
-415.89
-411.0
-288.0
-1384.49
-466.68
-1130.9
-947.68
0
-261.87
-376.56
-384.0
?
-1266.8
-365.89
-1047.67
-851.86
51.05
60.25
72.8
72.38
?
149.49
116.3
135.98
102.09
249.4
0
-12.09
142.2
230.1
0
16.3
163.4
160.95
205.0
110.88
237.6
0
-18.4
-1284.07
-3012.48
9.25
-1298.7
-1302.48
0
13.8
-1025.59
?
18.2
-1094.1
-1135.1
44.0
29.3
-217.57
?
210.0
-35.98
89.1
0
0.30
0
0.10
31.88
32.55
Substance
O (g)
O2 (g)
O3 (aq)
O3 (g)
P (white)
P (red)
PO43- (aq)
P4O10 (s)
PH3 (g)
HPO42- (aq)
H2PO4- (aq)
S (rhombic)
S (monoclinic)
UNIT 5
5
Substance
∆H˚f (kJ/mol)
∆G˚f (kJ/mol)
S˚ (J/K•mol)
SO2
SO3
SO32- (aq)
SO42- (aq)
H2S (g)
HSO3- (aq)
HSO4- (aq)
H2SO4 (l)
H2SO4 (aq)
SF6 (g)
-296.1
-395.2
-624.25
-907.5
-20.15
-627.98
-885.75
-811.3
-909.3
-1096.2
-300.4
-370.4
-497.06
-741.99
-33.0
-527.3
-752.87
-690.0
-744.5
-1105.3
248.5
256.2
43.5
17.15
205.64
132.38
126.86
156.9
20.1
291.8
Zn (s)
Zn2+ (aq)
ZnO (s)
ZnCl2 (s)
ZnS (s)
ZnSO4 (s)
0
-152.4
-348.0
-415.89
-202.9
-978.6
0
-147.2
-318.2
-369.26
-198.3
-871.6
41.6
-112.1
43.9
108.37
57.7
124.7
Substance
∆H˚f (kJ/mol)
∆G˚f (kJ/mol)
S˚ (J/K•mol)
-484.2
-389.45
159.83
-246.8
-153.55
198.74
226.6
209.2
200.8
49.04
124.5
172.8
-82
?
?
-276.98
-174.18
161.04
-84.7
-32.89
229.49
52.3
68.1
219.45
-1274.5
-910.56
212.1
-74.85
-50.8
186.19
-238.7
-166.3
126.78
-2221.7
-1544.3
360.24
CH3COOH (l)
(acetic acid)
CH3COCH3 (l)
(acetone)
C2H2 (g)
(acetylene)
C6H6 (l)
(benzene)
CH3Cl
(chloromethane)
C2H5OH (l)
(ethanol)
C2H6 (g)
(ethane)
C2H4 (g)
(ethylene)
C6H12O6 (s)
(glucose)
CH4 (g)
(methane)
CH3OH (l)
(methanol)
C12H22O11(s)
(sucrose)
UNIT 5
6
Hess’s Law Problems
1.
Given:
Find:
2.
S +
O2  SO3
ΔH = -395.2 kJ
2 SO2 + O2  2 SO3
ΔH = -198.2 kJ
S + O2  SO2
ΔH = ?????
Given the following data:
H2 + ½ O2  H2O + 286 kJ
N2O5 + H2O  2 HNO3 + 77kJ
N2 + 3 O2 + H2  2 HNO3 + 348 kJ
Calculate ΔHrxn for:
2 N2 + 5 O2  2 N2O5
UNIT 5
7
Enthalpy Calculation Problems
1. The equation for the reaction of hydrochloric acid with barium hydroxide solution is:
2 HCl + Ba(OH)2  BaCl2 + 2 H2O + 118 kJ
How much heat is produced if 34.5 g of HCl reacts with a stoichiometric amount of barium
hydroxide?
2. The combustion of acetylene (used in welding) is written as: C2H2 + O2  2 CO2 + H2O
Using heats of formation values, calculate ΔH° for this process.
3. Thermite is a generic term for a mixture of metal and metal oxide that is used to generate
tremendous heat. During the reaction, one metal is reduced and the other is oxidized. The
classic thermite mixture consists of iron (III) oxide and aluminum:
Fe2O3 + 2 Al  Al2O3 + 2 Fe
Using heats of formation, determine how much energy is released in this reaction.
UNIT 5
8
LAB: Bond Breaking and the Heat of Reaction
All chemical reactions take place with either an absorption or release of energy. Generally this energy is in
the form of heat, but in some processes it may take the form of mainly light, or a mixture of forms including some
mechanical energy such as sound. Whatever the case, the conclusion we could draw from this aspect of chemical
change is that energy is somehow stored in the chemicals and during a transformation there is sometimes excess
energy (in which case it will be released) or sometimes not enough energy (in which case it will be absorbed from
the surroundings).
The change in energy is due mainly to bond rearrangement during a reaction. Whether there is a net
absorption or release of energy depends on the number of bonds broken and formed as well as the strength of those
bonds. For example, in the combustion of methane:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
We could imagine the reaction taking place by all of the reactant bonds breaking to make individual atoms
and then these separated atoms later recombining to form the products. While it is unlikely that the actual chemical
process happens in this way, theory and experiment show that the net energy difference between reactants and
products is the same regardless of the pathway. Therefore, using bond energy values, we would have:
4(C-H) + 2(O=O) = 4(414 kJ) + 2(499 kJ) = energy to break bonds (+)
and
2(C=O) + 4(H-O) = 2(799 kJ) + 4(460 kJ) = energy to make bonds (-)
The sum of these two values is -784 kJ--in other words, 784 kJ is released when one mole of methane is
burned. The experimental heat of combustion is -803 kJ. So it is possible to estimate the energy change using a table
of experimentally determined bond energies, but such estimates are limited by a number of factors. The most
important limitation is the fact that the bond energies have been measured or calculated for gaseous species. In
addition, the energy required to break several similar bonds in a molecule is not the same. For example, the energy
needed to break one of the hydrogen-carbon bonds in methane is not the same as the energy needed to break the
second, third, or fourth one (although they are similar).
In this experiment you will be looking at a series of hydrocarbons from the alkane family. Alkanes have the
general formula CnH2n+2 where n is an integer. Thus methane (CH4) is the first member of the family. When you use
your bunsen burner in the lab you are burning methane and the heat given off is the heat of combustion (or, more
generally, the heat of reaction). Quantifying how much methane is burned in a given amount of time is difficult.
However, as the number of carbon atoms in alkanes increases, the compounds eventually become liquids at room
temperature and the mass of a compound burned is more easily measured. Three liquid alkanes will be used in this
experiment. They are hexane (C6H14), heptane (C7H16), and octane (C8H18). A structural formula for hexane is
shown below:
H
H
H
H
H
H
H
C
C
C
C
C
C
H
H
H
H
H
H
H
Heptane and octane are similar with longer carbon chains (the Greek prefixes contained in their names gives the
number of carbons in the molecule).
The combustion reactions for the alkanes are all similar. Carbon dioxide and water are products, as the
example for ethane below shows:
C2H6(g) +
7
2
O2(g)  2 CO2(g) + 3 H2O(ℓ)
While bond energy calculations give an estimate of heats of reaction it is also possible to measure the
energy change during a reaction in the laboratory. Such measurements are generally done in an insulated container
UNIT 5
9
(a calorimeter) that is otherwise open to the atmosphere so that the pressure remains constant during the reaction.
The heat that is absorbed or released under such conditions is known as Q. When this number is scaled to a molar
quantity it is called the enthalpy change of the reaction, ΔHrxn.
Reactions between liquids and solids are easily carried out in such a device or a reaction may be carried out
externally (as in this experiment) and the heat conducted to the calorimeter which is filled with water. A
thermometer can be used to follow heat flow during the reaction. If the temperature is observed to rise, this means
that heat is flowing from the system into the thermometer (and other surroundings). Such a reaction is described as
exothermic and the enthalpy change is given a negative (-) sign. If the temperature falls, this means that heat is being
absorbed from the surroundings into the system, indicating an endothermic reaction and making ΔHrxn positive (+).
In this experiment you will use a very simple constant pressure calorimeter which consists of a metal can
suspended over a small fuel burner. The can is surrounded by a larger can that helps to hold in some of the heat from
the fuel combustion and also limits drafts. Water is held in the smaller can and its temperature change during the
burning of the fuel (a combustion reaction) is used to determine the heat of reaction.
While this basic set-up is not insulated, it is possible to obtain good results from it if the experiment is
conducted in a consistent manner and if you know something about the heat loss characteristics of the arrangement.
The systematic error can be reduced by estimating the "calorimeter constant" or the heat capacity of the calorimeter
(i.e., the number of joules required to raise the temperature of the calorimeter by 1°C.). Actually, your measurement
will also include a factor for the heat that is lost from the calorimeter during the experiment. To determine the
calorimeter constant you will burn a fuel with a known heat of reaction (or heat of combustion in this case) and use
the known value to estimate the calorimeter constant. The following example illustrates this.
When 1.0 g of benzoic acid is burned in a calorimeter containing 2,950 g water, the temperature rises from
24.3 to 26.3oC. The heat of combustion for benzoic acid is -3,223 kJ/mol. What is the calorimeter constant
(i.e., heat capacity of the calorimeter)?
1) Heat (Q) = (mass x ΔT x specific heat) + (heat capacity of calorimeter x ΔT)
2) The heat of combustion is -3223 kJ/mol. In this case, we don't need the sign but we do need to scale
this value because we are only combusting one gram instead of one mole:
. x
. x
= 26 kJ
We must convert kJ to J to match units with the specific heat of water. Thus, the heat (Q) associated with
combusting 1.0 g of benzoic acid is 26000 J.
3) Heat (Q) = (mass x ΔT x specific heat) + (heat capacity of calorimeter x ΔT)
26000 J = (2950 g x 2.0oC x 4.184 J/goC) + (
C
x 2.0oC)
C = 660 J/oC
UNIT 5
10
In this experiment you will use the heat of reaction for the combustion of hexane (C6H14) as your reference
to determine the calorimeter constant. The measured value is -4.1 x 103 kJ/mol. Once you have used this value to
determine a calorimeter constant for the experimental set-up, you should be able to determine the heats of
combustion for heptane and octane. Converting the data you collect into molar heats of reaction (ΔHrxn) is illustrated
in the example below.
2.00 g of liquid pentane (C5H12) is burned and the heat produced is applied to a calorimeter containing
100.0 g water originally at 10.0°C. The final temperature of the water after reaction is 55.9°C. What is the
heat of reaction per mole of pentane? The calorimeter constant (heat capacity) is 1700 J/oC.
1) First we need to find out how much heat the reaction produces for the amount of pentane burned:
Heat (Q) = heat gained by water
+ heat gained/lost by calorimeter
Heat (Q) = (mass x ΔT x specific heat) + (heat capacity of calorimeter x ΔT)
Q = m s ΔT + C ΔT
Q = [100.0 g x 4.184 J/goC x (55.9-10.0oC)] + [1700 J/oC x (55.9-10.0oC)]
Q = 9.72 x 104 J
2) But this amount of heat is associated with 2.00 g of C5H12, not one mole! To determine the heat of
reaction per mole of pentane, we can use energy stoichiometry:
x
. x
.
.
= -3.5 x 103 kJ/mol pentane burned = Hrxn
(The negative sign was added to the final answer because the reaction is losing energy to the surroundings.)
Once you have determined the heats of combustion for heptane and octane then you can see if there is any
relationship between the number of carbon bonds and the energy released during combustion.
Preparing to experiment
You will be provided with the following materials:
1. a calorimeter set-up
2. a thermometer probe
3. three labeled fuel burners
4. hexane, heptane, octane
5. a 24-well plate
6. matches
7. a supply of water at approx. 10oC (use about 100 mL each time)
8. a 100 mL graduated cylinder
Design an experiment to determine the heat released by burning hexane, heptane and octane to raise the temperature
of a water sample by 20°C. For more information, see the Technique section below.
UNIT 5
11
Technique
1. Using the temperature probe
Using an electronic temperature probe is not quite the same as using an ordinary thermometer. The
response time is somewhat slower. For this first time here are some hints to help you get the most out of the
instrument.
a. Set up the LabQuest to record temperature.
b. Stir the water while heating so that the probe is not resting directly on the bottom of the can.
c. Continue to stir a little after extinguishing the flame to be sure the water temperature is the same throughout.
2. Using the calorimeter set-up
The calorimeters have been made so that the inner can will sit at the correct height when the fuel burner is
held in a 24-well plate. Be sure to use them that way. This will minimize heat escape. We also use one other trick to
help control heat loss. By starting with water below room temperature (which will slowly pick up heat from the
room) and ending with water about the same number of degrees above room temperature (which will slowly lose
heat to the room) we can effectively ignore that factor. Of course, not all of the heat from the flame goes into the
water, but that's what the determination of the calorimeter constant is all about. Incidentally, although it is better to
mass the water in the calorimeter rather than assume a density of 1.0 g/mL (since the graduated cylinder won't
deliver every last drop you measure out), in the interest of time, we will measure volume and assume the mass is the
same. Also, be sure to take a look at the bottom of the can after each heating. It needs to be clean and dry each time.
3. Using the fuel burners
All three alkanes evaporate pretty fast at room temperature. They also attack the plastic burner containers
slowly. For that reason, you will find the burners are empty. You can fill them from the bottles of the fuels provided
near the balances. They should only be filled about half-way. Be sure to empty them (into the correct beaker!!!!) at
the conclusion of the lab. Before lighting the burner, have all your preliminary measurements made and the mass of
the burner recorded. Inspect the wick and adjust it so that it is flush with the top of the tube and DOES NOT project
out of the tube. THIS IS VERY IMPORTANT. TOO MUCH WICK WILL PRODUCE A DANGEROUSLY
LARGE FLAME THAT COULD ENGULF THE ENTIRE BURNER. After heating, measure the mass of the
burner as soon as possible to minimize errors due to evaporation.
______________________________________________________
Analysis
1.
Use your data from the combustion of hexane to determine the calorimeter constant. [Note: Although the
example calculation on a previous page shows 1.0 g of compound used, you probably did NOT burn exactly 1.0
g of hexane in this experiment!!!]
2.
Calculate the heat or enthalpy of reaction (Hrxn) per mole of heptane and per mole of octane burned.
3.
The molar heat of combustion for pentane (C5H12) is -3.5 x 103 kJ/mol. Using this value for pentane, the value
given for hexane earlier in this lab, and the values you determined in question 2 for heptane and octane, graph
the molar heat of combustion vs. the number of C-C bonds in each compound. Does there appear to be a regular
relationship? If so, describe it and try to give a brief explanation based on bond energies. [Remember, energy is
absorbed to break bonds and released when bonds are formed.] As part of your answer, write balanced
equations for the combustions of hexane, heptane and octane, then compare the number of bonds formed in the
products for each reaction with the number of C-C bonds in the reactant]
[PS: Don’t be fooled!!! -5000 kJ is more heat than -4000 kJ]
UNIT 5
12
Name:___________________________ Per.:____
Date:_______________
The Determination of the Heat of Formation of Magnesium Oxide
Purpose: to determine the heat of formation of magnesium oxide indirectly through other
reactions involving magnesium, oxygen and magnesium oxide.
Method: I will measure the temperature changes for two reactions: magnesium + HCl and
magnesium oxide + HCl. For each reaction, I will use a measured volume of 2 M HCl in a
calorimeter, recording the initial temperature. I will then add either a 3.0 cm piece of Mg metal
or approx. 0.5 g of MgO while stirring. I will then measure the final temperature of the mixture.
From the temperature change, the given specific heat of the reaction mixture (s) and the given
calorimeter constant (C), I can calculate the amount of heat absorbed or released in each
reaction, using the following equation: Q = m s ΔT + C ΔT
I can then scale this heat for 1 mole of either Mg or MgO based on the quantities of these
limiting reagents. These two reaction heats can be used with the known heat of formation of
water (-286 kJ/mol) to determine the heat of formation of MgO using Hess’s Law:
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
MgO(s) + 2 HCl(aq)  MgCl2(aq) + H2O(ℓ)
H2(g) + ½ O2(g)  H2O(ℓ)
ΔH1
ΔH2
-286 kJ/mol
Data:
Table 1: Reaction of Mg + HCl
Table 2: Reaction of MgO + HCl
1.00 m Mg ribbon, g
(density of Mg ribbon)
length of Mg piece, cm
mass of MgO, g
cylinder + HCl, g
cylinder + HCl, g
cylinder – HCl, g
cylinder – HCl, g
initial temperature, oC
initial temperature, oC
final temperature, oC
final temperature, oC
UNIT 5
13
UNIT 5
14
LAB: The Determination of the Heat of Formation of Magnesium Oxide
Chemical reactions proceed with the evolution or absorption of heat. This heat flow represents differences
in chemical energy associated with the arrangement of atoms into molecules or ions. It includes the electrical
potential energy arising from attractions between electrons and nuclei in atoms or ions for each other. Thus when a
piece of magnesium ribbon is burned in air to yield magnesium oxide, a large quantity of heat is released to the
surroundings:
Mg(s) + ½ O2(g)  MgO(s) + heat
Since heat is evolved in this reaction, we know that the sum of the chemical energies of the amounts of
elemental magnesium and oxygen combining is greater than the chemical energy of the amount of magnesium oxide
formed.
When the reaction takes place at constant pressure, the heat flow, Q, is called the change in ENTHALPY of
the system and for MOLAR quantities is given the symbol ΔH. The absolute value of the enthalpy, H, for a
substance is indeterminate, but ΔH for a particular process is directly measurable, at least in principle, and is a
quantity of fundamental interest in thermochemistry and thermodynamics.
When one mole of a substance is formed from its constituent elements in their standard states at standard
conditions, ΔH for the reaction is called the standard heat of formation and given the symbol ΔHof . In this
experiment you will try to find ΔHof for magnesium oxide, i.e., for the reaction:
Mg(s) + ½ O2(g)  MgO(s)
You are faced with a practical difficulty, however. To cause magnesium and oxygen to react they must be
heated. The reaction then proceeds in an uncontrolled and highly exothermic fashion. The only device that can deal
with this reaction directly is a bomb calorimeter…and we don't own one of those. However, so long as the system
starts at and returns to standard conditions, the path it follows is not important. Any series of reactions that begins
with Mg and O2 and ends with MgO will involve the same overall change in enthalpy. Thus you can use Hess's
Law and a number of reactions which are more easily dealt with to find the information about the reaction of
magnesium and oxygen. Consider the following reactions:
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
(1)
MgO(s) + 2 HCl(aq)  MgCl2(aq) + H2O(ℓ)
(2)
H2(g) + ½ O2(g)  H2O(ℓ)
(3)
For the first two reactions, the heats of reaction are easily found in a simple calorimeter. For the third
reaction, the heat of formation of liquid water, the enthalpy change is well known (-286 kJ/mol). By reversing
equation (2) and adding all three we obtain:
Mg(s) + ½ O2(g)  MgO(s)
(4)
The enthalpy change for this reaction (ΔHof), the heat of formation for magnesium oxide, is found by a
similar reversal and addition of the enthalpies of the three reactions.
The constant-pressure calorimeter used in this experiment consists of a specially treated block of expanded
polystyrene foam (commonly called “styrofoam”, but not the same stuff as the rough and porous material often used
in crafts). The block has a hole drilled in it with a practical volume of about 15 mL and is coated to help prevent it
from absorbing liquid. Although the calorimeter is fairly well insulated some heat will escape or enter through hole
in the lid provided for the thermometer, and some heat will be absorbed into or from the calorimeter contents by the
calorimeter itself. For the purposes of this experiment we will ignore the former random error. But the systematic
error introduced by ignoring the calorimeter itself can be reduced by using the calorimeter constant in heat
calculations. The average calorimeter constant for this set-up has been determined to be 5.6 J/oC.
UNIT 5
15
Preparing to experiment
You will be provided with the following materials:
1. 3.0 cm of Mg ribbon [The mass in grams per meter will be given!]
2. 2.0 M HCl - use about 15 mL each time [see Technique section]
3. a calorimeter, lid and stirring bar
4. a magnetic stirrer
5. MgO - use about 0.5 g [see Technique section]
6. a thermometer probe
Design an experiment to determine the heat of reaction when a measured amount of
Mg reacts with HCl (the amounts suggested make Mg the limiting reagent).
Design an experiment to determine the heat of reaction when a measured amount of MgO reacts with HCl (the
amounts suggested make MgO the limiting reagent).
Pre-lab take-home quiz
Answer these questions on a separate sheet of paper to be turned in on the day you do this experiment.
1.
Most Mg ribbon is uniform enough that the mass of a small measured length is proportional to the mass of a
larger measured length. Assume that the mass of 1.00 meter of Mg ribbon is 0.4368 g. Show that using 3.0 cm
of this ribbon and 15 mL of 2.0 M HCl makes Mg the limiting reagent in the reaction between the two
substances.
2.
Show how the three reactions given in the introduction actually add up to give the formation reaction for MgO
(i.e., set up the three equations and show how they add up, canceling like terms, etc. using Hess’s Law.)
Technique
1. Transfer errors
In calculating the heat released by the reactions it is important to know the mass of the reaction mixture.
Because we are using small amounts in this experiment we need to try and minimize the error that would normally
not be very significant when, for example, liquid is added from a graduated cylinder.
In the Carbonate Project you used a technique of massing the amount of liquid added to a reaction by the
difference of the graduated cylinder full of the liquid and the cylinder after the liquid had been poured out. That
would be appropriate in this experiment once again.
The MgO is a very light powder and tends to stick to glass to some extent so it too is a problem if you plan
on measuring some into a beaker and dumping it into the calorimeter. You could mass by difference but there is the
additional problem of adding the powder to a calorimeter full of HCl. The reaction begins immediately at the surface
and heat is lost (spray too) before the lid is replaced. One way to reduce this error is to measure the MgO directly
into the calorimeter and then add the HCl quickly. For this to work it is important that the calorimeter is dry.
2. Drying the calorimeter
These blocks are fairly soft and you can damage the surface coating if you are careless handling them.
Drying the inside with our normal paper toweling requires a gentle touch. A final drying with a softer Kimwipe is a
good idea.
UNIT 5
16
3. Dealing with the thermometer probe
You need to suspend the probe so that the tip is in the solution but will not hit the stirring bar. One way to
do this is to use your ring and ring stand. The probe cable can be draped over the ring and the probe will dangle over
the calorimeter. A little adjustment to line things up, and it’s done!
The chemicals
Magnesium oxide occurs in nature as the mineral periclase. It is a white, very fine, odorless powder. It gradually
picks up CO2 and water from the air, combining with the water to form magnesium hydroxide (milk of magnesia).
Magnesium oxide is used the manufacture of refractory crucibles and bricks as well as in medical applications as a
laxative and antacid.
Analysis
1.
Use your data to determine the heat of reaction per mole of Mg for the reaction between Mg and HCl. You did
not use one mole of Mg, but enthalpy changes are proportional to moles. Don't forget to include the calorimeter
constant in the calculations. [The specific heat of the reaction mixture is 3.75 J/goC.]
2.
Use your data to determine the heat of reaction per mole of MgO for the reaction between MgO and HCl. You
did not use one mole of MgO, but enthalpy changes are proportional to moles. Don't forget to include the
calorimeter constant in the calculations. [The specific heat of the reaction mixture is 3.75 J/goC.]
3.
Use the molar enthalpies from (1) and (2) above, and the enthalpy value given earlier for the formation of water
to determine the heat of formation for MgO (ΔHof) as you demonstrated in your answer to question #2 in the
pre-lab quiz.
4.
With the final result, the question arises (as it does in any experiment) as to its accuracy. Using values obtained
from handbooks, the following heats of reaction have been calculated:
Reaction (1)
-462 kJ/mol Mg
Reaction (2)
-146 kJ/mol MgO
Reaction (4)
-602 kJ/mol MgO
For each, determine the % error* in your experimental results and briefly discuss any significant deviations with
specific reference to things you did and would not do again, etc.
*
UNIT 5
experimental - actual
actual
x 100
17
UNIT 5
18
Unit 5 Sample Test
The test will be similar in format to previous tests. Recall that you still have a choice of two out
of four reactions to complete and write balanced net-ionic equations for. You will also need to
identify the type of reaction. This time the first four solubility rules will be missing from the test.
_____1.
Ionic bonds form between
a. non-metals with high electron affinity and metals with low ionization energy
b. non-metals with low electron affinity and metals with high ionization energy
c. metals with high electron affinity and non-metals with high ionization energy
d. metals with low electron affinity and non-metals with low ionization energy
_____2.
Consider the reaction below:
1.7 kJ + NaCl(s) + H2O(ℓ)  Na+(aq) + Cl-(aq)
Which of the following combinations for ΔH and ΔS is correct?
a. +,-
_____3.
b. -,+
c. -,-
d. +,+
The change in enthalpy, ΔH, is equal to the heat gained or lost by the system when
a. pressure is constant
b. volume is constant
c. temperature is constant
d. none of these
For the following question, compare the assertion (top) and the reason (bottom) and select the
appropriate answer from the options given below:
a. both are true and reason fits the assertion
b. both are true but reason does not explain assertion
c. only the assertion is true
d. only the reason is a true statement
e. both are false
_____4.
The vaporization of water at room temperature is not a “spontaneous” process
BECAUSE
the temperature is such that TΔS is less than ΔH.
_____5.
Which of the following reactions is likely to have the greatest increase in entropy?
a. 2 Na(s) + Cl2(g)  2 NaCl(s)
b. NH4Cl(s)  NH3(g) + HCl(g)
c. 2 Mg(s) + O2(g)  2 MgO(s)
d. NaCl(s) + H2O(ℓ)  Na+(aq) + Cl-(aq)
UNIT 5
19
The next section consists of typical problems like those that might be found on the test.
6. a. Write a balanced molecular equation for the burning of 1 mole of C2H4 to form carbon
dioxide and liquid water, and calculate ΔHorxn using values from your reference pages.
b. If 15.4 grams of C2H4 is burned instead of the 1 mole, how many kJ will be released?
c. The entropy change (ΔSorxn) for the reaction in (a) is -267.45 J/mol·K. Calculate ΔGorxn
for this reaction at 298 K.
d. Explain how this reaction can be spontaneous at 298 K even though ΔSsys is negative.
7. A sample of carbon with a mass of 0.463 g was placed in a calorimeter containing 2500 g of
water. As a result of burning the carbon, the temperature of the water rose from 22.54°C to
23.82°C. The calorimeter constant is 1460 J/oC.
a. Write a reaction for the combination of the carbon with elemental oxygen.
b. What is ΔHrxn? Include a correct sign on your final answer.
UNIT 5
20
8. Calculate ΔHorxn for the reaction shown below using bond enthalpies and heats of formation.
Discuss briefly the reason(s) for the differences in the values you obtain for ΔHorxn using
each method.
9. Given the following reactions:
2 P (s) + 3 Cl2 (g) → 2 PCl3 (ℓ)
ΔHo = -635.1 kJ
PCl3 (ℓ) + Cl2 (g) → PCl5 (s)
ΔHo = -137.3 kJ
Calculate the heat of formation, ΔHof, for solid PCl5 under the same conditions. [Hint: You
must first write the reaction for the formation of PCl5.]
10. Indicate whether each statement is true (T) or false (F). If the statement is false, correct it.
_____a. Reactions for which energy must be supplied from outside the system to carry out
the processes are spontaneous.
_____b. Spontaneous processes always occur very rapidly.
_____c. The entropy of a substance in the gas phase is always greater than the entropy of the
same substance in the liquid phase.
_____d. According to the second law of thermodynamics, when ΔSsys > 0, the process is
spontaneous.
_____e. A reaction that is spontaneous in one direction is non-spontaneous in the reverse
direction under the same conditions.
UNIT 5
21
The next section consists of representative reactions to complete and write balanced net-ionic
equations for. Note that some reactions do not occur in aqueous solution and thus molecular
equations are all that would be needed. Each student is expected to choose two from this section.
In addition to a periodic table and a partial set of solubility rules, the test will include a copy of
the activity series.
11. For each of the following, complete the word equation, write a balanced net-ionic reaction
and tell what type of reaction it is. (b, c and d occur in water solution)
a. carbon monoxide + oxygen → (one product)_____________________________________
___________________________________________________________________________
___________________________________________________________________________
type:______________________
b. sodium chloride + lead(II) nitrate → ___________________________________________
___________________________________________________________________________
___________________________________________________________________________
type:___________________
c. potassium metal + water → ________________________________________________
___________________________________________________________________________
___________________________________________________________________________
type:_____________________
*d. dichromate ion reacts with ethanol (C2H5OH) in an acidic solution; chromium(III) ions and
carbon dioxide are among the products.
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
The final section of the test will consist of one essay question selected from the following topics:
--effects of experimenter errors in calorimetry
--enthalpy, entropy and spontaneity: how are they connected?
UNIT 5
22