Download LaurenHill Chemistry 534

LaurenHill Chemistry 534
5.
Energy and Enthalpy
A.
What is Energy?
Energy is the capacity to do work. Although the
definition may not seem to do justice to a poet's concept of
energy, think of any of its forms, and it involves work. For
instance, visible light does work on any material that absorbs it,
causing molecules to vibrate and electrons to jump levels.
Example
Give another example of how a form of energy is a capacity to do work.
Energy is neither created nor destroyed; it is merely transformed from one form
to another. For example in moving a muscle, chemical energy is transformed into
mechanical energy; if you move backwards, you can actually trace every little body
action to the sun.
How?
B.
Potential Energy and Enthalpy
To understand the difference between two types of reactions (exothermic and
endothermic), we need to explore a couple of other concepts. In addition to kinetic
energy (vibrational, rotational and translational motion), molecules also have potential
energy. Potential energy in a chemical context is energy due to composition. It is
associated with the coulombic force within atoms and molecules (intramolecular
bonds); with the attractions between different molecules (intermolecular bonds).
In water for example, potential energy is associated with bonds between
oxygen and the two hydrogen atoms in each molecule, and also with the oxygen atom
of one molecule and one of the hydrogen atoms of another molecule.
Intramolecular bonds
H
H
O
H
H
Intermolecular bonds
O
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II. Thermochemistry
Example1
Show the intermolecular and intramolecular bonds in HCl.
Example 2
a)Imagine two pieces of plasticine bonded together. If they represent a diatomic
molecule, what kind of bond is represented? b)What holds the real pieces
together?
Example 3
How do bonds between protein molecules play a role in the cooking of egg
white? ( see http://www.sumanasinc.com/webcontent/animations/content/proteinstructure.html )
Enthalpy Definition: H is the total heat content of a substance at constant pressure.1
C.
Exothermic reactions
If in a reaction molecule A becomes molecules B and C, and if molecule A has more energy
that both B and C combined, then the excess energy will be released into the environment. The
environment becomes hotter; we have an exothermic reaction:
A

Reactant
Examples of exothermic reactions:




B + C + energy
products
Condensation of water
C + O2  CO2 + 250kJ
Adding an alkali metal to water (2 Na + 2 H2O  2 NaOH + H2 + energy)
All combustion reactions (fires)
1
Enthalpy is also defined as the internal energy of the system plus the product of pressure difference and
volume(work). But if pressure is constant, then it is simply equal to its internal energy.
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LaurenHill Chemistry 534
Example
Think of everyday reactions which release heat (exothermic).
On a graph, exothermic reactions are represented as follows:
Energy
reactants
products
Progress of reaction
If we examine the graph more closely, we will notice that exothermic reactions have a negative
change in enthalpy. A change in enthalpy, H, is defined as the enthalpy of products – heat of
reactants:
H = Hp - Hr
What is that little
hill labeled, Ae? Ae
= activation energy.
This is the energy
that reactants must
absorb in order to
form products, even
if the products will
not need the energy
to store within their bonds. So Ae = Hmaximum - Hreactants
Example 1
Find H if the reactants have 120 kJ, and the products have 40 kJ of enthalpy.
Example 2
If a reaction has a negative value for H, what will happen to the temperature
of its environment?
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II. Thermochemistry
Example 3
When chlorine and gently heated sodium are placed in the same beaker there is
an incredible amount of light and heat released as they react to form table salt, NaCl. Where
does that energy come from?
D.
Endothermic Reactions
If substance A must take energy away from the environment in order to form product D, then
the reaction is said to be endothermic, and the victimized environment will feel colder after the
reaction.
H = (+) for endothermic reactions and their profile looks like the following:
Examples of endothermic reactions:



Melting of ice absorbs energy
Dissolving ammonium nitrate in water (the essence of commercial
cold packs)
NH4 NO3(s) + energy NH4 NO3(aq)
Example 1
Think of everyday reactions that take heat away from their environment.
(endothermic)
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LaurenHill Chemistry 534
Example 2
If a reaction has a positive value for H, what will happen to the temperature of
its environment?
Check out these next two pages for an analogy between physics and chemistry so that you
gain a better understanding of potential energy.
Potential Energy from Physics Point of View
Energy
Work is done
against gravity
to increase the
height of the
object.
Ep
= (mg)h
Work = ( Fg )d
After work is done against the force of gravity(Fg) to raise the object to a greater height, the
energy is conserved as potential energy .
If the box falls it will release the energy in the form of kinetic energy, which can do ”work “ on
the floor, a spring, or whatever your fancy.
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LaurenHill Chemistry 534
Potential Energy from the Chemistry Point of View but with Physics also in Mind
I2(g) + energy = 2 I(g)
I.
Energy
.
I
..
II
Work is done against coulombic attraction
to increase the distance between what were
shared electrons(-) and the nucleus(+).
Work = ( Fc )d
Which becomes chemical potential energy
When iodine atoms are sharing two electrons in the I2 molecule, the shared electrons are much
closer to each other’s nucleus. The separation distance, d, is small, just like the height in the
physics analogy.
To break up the bond requires energy because you are separating the shared electrons as they
become two separate atoms. “d” has increased and so has the potential energy. We have an
endothermic reaction.
In every form of chem. potential energy (intraatomic, intramolecular or intermolecular, a similar
idea is at work.)
Exercises (the last few exercises are from old exams)
1. The Romantic poet William Blake defined energy as eternal
delight. But from a more practical perspective, just what is
energy?
2. The energy associated with physical or chemical changes may
appear in different forms such as heat or electricity.
List at least two other forms of energy.
3. A 5.0 gram mass is moving at 1000 m/s. Calculate its kinetic
energy without forgetting to convert mass to kilograms.
4.
What is the difference between potential and kinetic energy?
5.
Give 3 examples of potential energy.
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II. Thermochemistry
6.
The Law of Conservation of Energy states that energy can be transformed from one form to
the other, yet it cannot be created or destroyed. What other Law important to balancing
equations does this remind you of?
7.
a. A solar cell absorbs light. Into what form of energy is light transformed?
b. If the electricity produced is allowed to flow into a toaster, what two
forms of energy is electricity turned into?
8.
a. Discuss the three types of kinetic energy associated with molecules.
b. Describe what happens to the molecules or atoms of a solid as its temperature rises to
its melting point.
c. Repeat (b) for the condensation of water.
9.
a. Why do molecules have potential energy?
b. What is the difference between intermolecular and intramolecular bonds?
c. List 3 everyday examples where the potential energy of a chemical substance is
released.
10. What happens to the energy of molecules that react chemically?
11. What is enthalpy?
12. What is an endothermic reaction? Exothermic?
13.
Record and complete the following:
H of reactants (kJ)
H of products (kJ)
44
288
44
555
222
33
73
88
41
14.
Enthalpy Change,
(kJ)
Exo or Endo?
-55
444
255
222
-88
55
38
a.
Draw an energy-versus-progress-of-reaction profile, using numbers from the
above chart's first row.
b.
Repeat for second-row data.
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LaurenHill Chemistry 534
15.
A student dissolved 1 g of each of four substances in water in a laboratory experiment.
The following table illustrates the change in temperature after the solids dissolved.
SUBSTANCES
Tinitial
Tfinal
NH4Cl
23°C
19°C
NaOH
23°C
60°C
NaCl
23°C
23°C
Drano
23°C
60°C
Which of the above represent exothermic reactions?
16.
Endothermic? Or Exothermic?
1. H2 O(g)  H2 O(l) + e nergy
2. H2 +
1
O2  H 2 O
2
H = - 241.8 kJ
1
3. H2 SO4(l) + h eat  H2(g) + S8(s) + 2 O2(g)
8
4. N2(g) + O2(g)  2 NO(g) H = + 180.8 kJ
5. C(s) + O2(g)  CO2(g) + 394 kJ
17.
A lighter has a reservoir containing a flammable gas under pressure.
When the lighter is ignited, two things happen:
o
o
gaseous butane escapes through an opening;
friction from the flint creates heat and sparks
The combustion of butane is exothermic.
Associate the expressions on the right with the numbers on the energy diagram on the left.
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II. Thermochemistry
Energy Diagram
Expressions
E (kJ)
}
}
a.
The energy contained in butane when the
lighter is opened
b.
The energy associated with the butane
heated by the sparks
c.
The energy liberated by the combustion of
butane
d.
The energy contained in the products
formed by the combustion of butane (CO2(g)
and H2O(g)).
3
1
4
2
5
Reaction Proceeds
18.
Which of the following are endothermic changes?
a.
b.
c.
d.
e.
f.
Melting ice
A burning candle
Dew forming on a lawn
Moth balls undergoing sublimation
Iron rusting
Water decomposing by electrolysis
6.
Hess' Law
Hess' Law states that the enthalpy of a reaction is independent of
whether the reaction occurs in one or several steps. This allows us to
algebraically add equations and their accompanying H's to obtain the
H for the desired or target equation.
Keep the following rules in mind:
Germain Henri Hess
1. If an equation is multiplied or divided by a number, that factor also applies to H.
2. If an equation is reversed, then the sign of H changes.
3. Remember: there are different enthalpy changes associated with different states of
matter. Do not, for example interchange H2O(l) with H2O(s).
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LaurenHill Chemistry 534
Example 1
Carbon disulfide is a very flammable solvent. It burns according to the
following equation:
CS2(l) + 3 O2(g)  CO2(g) + 2 SO2(g)
Calculate H for the above reaction using the following data:
(1) C(s) + 2 S(s)  CS2(l)
(2) C(s) + O2(g)  CO2(g)
(3) S(s) + O2(g)  SO2(g)
H = 88 kJ
H = -394 kJ
H = -297 kJ
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II. Thermochemistry
Example 2
Given:
Al2O3(s) + 1676 kJ  2 Al(s) +
H2O(g)  H2(g) +
1
O2(g)
2
3
O2(g)
2
H = + 242
kJ
mol
What is the energy change associated with the following reaction in terms of
kJ
?
mol Al
Al2O3(s) + 3 H2(g)  2 Al(s) + 3 H2O(g)
Example 3
Use the table from HW question 2 to find the heat of combustion for
ethane (C2H6).
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LaurenHill Chemistry 534
Exercises
1.
Methane (CH4) burns in the presence of oxygen according to this balanced
equation :
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
Below are several useful equations:
H2(g) + 1/2 O2(g)  H2O(g) + 242 kJ
C(s) + O2(g)  CO2(g)
H = -394 kJ/mol
C(s) + 2 H2(g)  CH4(g) + 75 kJ
What is the heat of combustion for 1 mole of methane?
2.
The molar heat of reactions of various elements are listed below:
Elements
H2(g) +
1
O2(g)
2

1
S8(s) + O2(g)
8
1
H2(g) + S8(s) + 2O2(g) 
8
1
C(s) + O2(g)
2
C(s) + O2(g)
C(s) + 2 H2(g)
2 C(s) + 3 H2(g)
3 C(s) + 4 H2(g)
H (kJ)
Formula
Name
H2 O(g)
water vapour
-241.8
sulfur dioxide
-296.9
H2 SO4(l)
sulfuric acid
-811.4

CO(g)
carbon monoxide
-110.5


CO2(g)
CH4(g)
carbon dioxide
methane
-393.5
- 74.8


C2 H6(g)
C3 H8(g)
ethane
propane
- 84.7
-103.8

SO2(g)
1
O2(g)  CO(g)
2
a.
Calculate H for
CH4(g) +
b.
Calculate H for
C3 H8(g) + 5 O2(g)  CO2(g) + 4 H2 O(g)
+ 2 H 2(g)
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II. Thermochemistry
3.
Calculate the heat of combustion for ammonia, NH3(g).
12 NH3(g) + 21 O2(g) 
8 HNO3(aq)
Given:
NH3(g) +
4.
a.
1.25 O2(g)

+
14 H2O(l)
+ 4 NO(g)
NO(g)
+ 1.5 H2O(l)
H = -293 kJ
NO2(g) + (1/3) H2O(l) 
(2/3) HNO3(aq) + (1/3) NO(g) H = -45 kJ
NO(g) + 0.5 O2(g) + 59 kJ

NO2(g)
Find the heat of reaction for the combustion of nitric oxide:

2 NO(g) + O2(g)
2 NO2(g)
The following experimental data is available to you:
b.
5.
0.5 N2(g)
+
0.5 O2(g)
0.5 N2(g)
+
O2(g)


NO2 (g)
NO (g) + 93 kJ
H = - 34 kJ
If you had used the data from #3, could you have gotten the same answer
in one step? How?
Diborane, B2H6, is highly reactive and was once considered as a possible rocket
fuel for the U.S. space program. Calculate H for the synthesis of diborane from
its constituent elements:
2 B(s) + 3H2(g) 
2 B(s) + 1.5 O2(g)
B2H6(l)
Reaction
 B2O3(s)
H ( kJ )
-1273
B2H6(l) + 3 O2(g)  B2O3(s) + 3 H2O(g)
-2035
H2(g) + 0.5 O2(g) 
-286
H2O(g) H2O(l)
H2O(l)
-44
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LaurenHill Chemistry 534
Estimating H From Bond Energies
It takes energy to break old bonds but energy is released when new bonds are formed. If
we could do the chemical accounting and sum up what’s invested and what the returns
are, we could estimate the net result or the H for a reaction.
Important reminder:
kJ
Energy invested to break bonds; Hbb = (+)
Energy released when bonds are formed; Hbf = (-)
7.
Hbb
Example 1: Estimate the  per mole offor the following
reaction:
H2 + Br2  2 HBr
H
Br
H
Br
Br
H
Br
H
hydrogen bromide
bromine
hydrogen
progress
H
Hbf
Rxn progress
Hbb
Hbf
H =HbbHbf
Table of Bond Energies to consult:
Average bond energies, kJ/mol:
H—H
436
H—F
570
H—Cl
432
H—Br
366
H—I
298
C—C
347
C=C
619
C≡C
812
C—N
293
C=N
515
C≡N
891
C—H
C—O
C=O
C≡O
C—F
C—Cl
C—Br
C—I
N—H
N—O
N=O
414
335
745
1075
485
326
285
239
389
175
590
N—N
N=N
N≡N
N—F
N—Cl
O—O
O=O
O—H
F—F
F—Cl
159
418
941
270
201
138
494
464
159
256
P—H
P—Cl
S—H
Cl—Cl
Br—Br
Br—Cl
I—I
318
326
339
243
193
218
151
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II. Thermochemistry
Example 2: Estimate the for the following reaction. Again, consult the bond
energies table on the previous page:
CH4 + 2 O2  CO2 + 2 H2O
Lewis Structures
Hbb
Hbf
H
=HbbHbf
Example 3
Show an energy- reaction profile (energy diagram) for an endothermic
reaction in terms of Hbband Hbf
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LaurenHill Chemistry 534
Exercises
1.
Estimate H for each of the following using the given table of bond energies:
a)
2H2
+ O2

b)
C2H6
+
3.5 O2 2 CO2 + 3 H2O
c)
2 CH3OH + 3 O2  2 CO2 + 4 H2O
Average bond energies, kJ/mol:
H—H
436
H—F
570
H—Cl
432
H—Br
366
H—I
298
C—C
347
C=C
619
C≡C
812
C—N
293
C=N
515
C≡N
891
2.
2 H2O
C—H
C—O
C=O
C≡O
C—F
C—Cl
C—Br
C—I
N—H
N—O
N=O
414
335
745
1075
485
326
285
239
389
175
590
N—N
N=N
N≡N
N—F
N—Cl
O—O
O=O
O—H
F—F
F—Cl
159
418
941
270
201
138
494
464
159
256
P—H
P—Cl
S—H
Cl—Cl
Br—Br
Br—Cl
I—I
318
326
339
243
193
218
151
a) Without consulting a table of bond energies, find the bond energy of H--Cl on a per
mole basis if H2’s and Cl2’s bond energies are 436 kJ/mole and 243 kJ/mole,
respectively, and the H for the following reaction is
-185 kJ/mole of H2
H2 + Cl22 HCl
b)
Convert H2’s bond energy to kJ/g
c)
Show an energy- reaction profile (energy
diagram) for
H2 + Cl22HCl in terms of Hbband Hbf
3.
Why is H---Cl’s bond energy greater than that of H--Br?
Hint:
4.
periodic trends.
a) Explain why the first 3 steps of the salt-producing
reaction are endothermic.
b) Why are the last two steps are exothermic?
c) Use algebra to show that the overall reaction is
indeed Na(s) + ½ Cl2(g) NaCl(s).
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II. Thermochemistry
8.
Calorimetry and Molar Enthalpy
The purpose of calorimetry is to use an instrument
known as a calorimeter to determine the enthalpy
of a substance undergoing chemical change. In a
calorimeter known as a bomb calorimeter, it is the
enthalpy of combustion that is measured. This is
how the caloric content of foods is determined. In
both cases, since the heat absorbed or released is
proportional to the amount of reactant used, molar
enthalpy = H/n is a more meaningful and
characteristic quantity.
In a bomb calorimeter, the actual chamber holding
the sample is known as a “bomb”. After opening
its lid, we place a weighed sample in a cup at the
bottom of the bomb. We seal it, and through a
valve, we deliver O2, saturating the bomb to get it
ready for ignition. We secure the bomb within a
calorimeter bucket that is filled with water (the water is the environment which will
absorb the heat of combustion). A stirrer keeps the temperature of the water evenly
distributed. A thermometer allows us to measure the initial temperature; the ignition wire
connected to a high voltage source initiates the explosion; heat is released, and we
measure the maximum temperature attained.
Q = mc T
m = mass of the water in the calorimeter in grams (because of c ‘s units; see below), or
the mass of whatever substance is acting as the environment. In reality we should assume
that the material part of the calorimeter also absorbs heat. But in this course, we usually
ignore that part.
c = specific heat of water or whatever is acting as the environment.
Note c is usually expressed in
, so that Q is in Joules
T = Tf – Ti.
To get H, remember:
H = - Q.
( see "I'm so hot and she's so cold”, on page 43")
We convert J to kJ by  1000, and then
Molar enthalpy = H/n.
n = number of moles of reactant. So we convert the
carefully measured mass in to moles by dividing by molar
mass.
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LaurenHill Chemistry 534
In molar heat of neutralization problems, n = CV, where
C = concentration in “M” = moles/L.
V = volume in liters.
In-Class Examples
1.
9.0 grams of charcoal (C) were completely consumed in a bomb calorimeter. If
we assume that the 2.0 L of water absorbed all of the heat released by the
charcoal, and if the temperature of the water increased from 20.25 to 56.04oC,
what is the molar enthalpy of carbon?
2.
CS2, a very flammable liquid, has a molar enthalpy of -1028 kJ/mole. What do
you expect aluminum's final temperature to be if 1.0 kg of Al is initially at 20.0 C,
and it absorbs all the heat from the following sample of CS2:
mass of CS2 before burning: 22.6 g
mass of CS2 after burning: 11.6 g
specific heat capacity of Al: 0.900 J/[g C]
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II. Thermochemistry
3.
300.0 mL of 0.20 M aqueous KOH neutralizes 150 mL of aqueous 0.20 M H2SO4.
We go from an average initial temperature of 22.3 C to a maximum of 29.2 oC.
Calculate the molar heat(enthalpy) of neutralization of KOH. Assume that the
amount of water created in the neutralization is negligible and that the specific
heat and density of all aqueous solutions is 4.19J/(goC) and 1.0 g/ml, respectively.
4.
Find the final temperature of the following mixture:
400.0 g of Cu initially at 99.0 oC
25.0 L of water initially at 10.0 oC
c for Cu = 0.39 J/[g C]
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LaurenHill Chemistry 534
Example:
Relate the "He's so hot and she's so cold" to why Q = -H.
Throughout the explanation, imagine a man with hot hands and a woman with cold
hands.
1st Point of View: Man is reaction; woman is environment. Man
releases or loses heat; he represents
an exothermic reaction. H for man is (-). Woman feels heat; she
absorbs it, and her temperature increases. Her Q = (+) . Recall that Q
= mc , where  = Tf - Ti.
Note H = -Q.
2nd Point of View: Woman is reaction; man is environment. Woman
takes heat from man. She is an endothermic reaction. She steals or
absorbs heat from the man. H for woman is (+).
Man feels a cold hand; he loses heat to the woman, so his temperature drops. His Q = (-).
Again H = -Q; the signs are again opposite of one another.
Chemical Case.
Suppose that NaOH(aq) is neutralizing HCl(aq). Both are dilute solutions in water. When
they react they release heat into the surrounding water, just like the man. Their H = (-).
But you record an increase in temperature because you are inserting a thermometer into
the water. The water is like the woman in the previous analogy. So for water, Q = (+).
If NaNO3 dissolves in water, the sodium nitrate absorbs or steals heat from the water. The
H for the reacting nitrate is (+), just like the woman's when she is regarded as the
reactant. The water is now like the man; it experiences a drop in temperature and its Q is
(-).
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II. Thermochemistry
Exercises (questions slightly modified from old exams)
1.
Calculate the heat released by the combustion of one mole of C4H10(g) based on
the data below:
initial mass of butane
remaining mass of butane
water's initial temperature
water's final temperature
volume of water in calorimeter
7.25 g
1.25 g
25.0 oC
60.7 oC
2.0 L
2.
Find the molar heat of KOH(s) if dissolving 4.0 g of this base in 200.0 mL of
water caused the water in the calorimeter raised from 25.0 to 31.5 oC.
3.
When 200 mL of 0.50 moles/L NaOH solution completely neutralizes 200 ml of
HCl, the temperature of the solution rises by 6.6 oC. Calculate the molar heat of
neutralization of NaOH, assuming that the resulting solution has the same density
and specific heat as those of water.
4.
In an experiment, a piece of iron, Fe, with a mass of 4.0 grams was heated from
22°C to 42°C.
During the heating process, the piece of iron absorbed 5.04  101 kJ/mol. What
is the specific heat capacity of this piece of iron?
5.
Uncle Ivan prefers his pre-dinner drink to be about 15°C. What mass of tap water
(at 11°C) must he add to his 30.0 mL of alcohol at 22°C so that his drink can be at
his preferred temperature (15°C)?
(Assume the alcohol has the same density and specific heat capacity as ethanol =
J
; density of ethanol = 0.79g/ml)
2.45
g  C
6.
A piece of metal with a mass of 32.6 grams at a temperature of 2.00  102°C is
dropped into a calorimeter containing 1.00  102 mL of water at 25.0°C.
The specific heat capacity of the metal is 0.448
J
.
g  C
Assuming complete heat transfer between the water and the metal, what will be
the maximum temperature of the metal-water system in the calorimeter?
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LaurenHill Chemistry 534
7.
A student was asked to check the purity of the paraffin, C25H52, in a candle.
Knowing that the molar heat of combustion of pure paraffin is 15 300 kJ/mol, she
conducted an experiment to determine the molar heat of combustion of the
paraffin in the candle. She used a calorimeter and made the following
observations:
Mass of the candle at the beginning of the experiment
Mass of the candle at the end of the experiment
Initial temperature of the water in the calorimeter
Final temperature of the water in the calorimeter
Volume of water in the calorimeter
28.52 g
24.29 g
20.0°C
42.0°C
2000. mL
Note: All of the heat released during the combustion of the paraffin was absorbed
by the water in the calorimeter.(question on next page……)
Can the student conclude that the paraffin in the candle is relatively pure, given
that the paraffin in the candle is considered pure if the difference between its
molar heat of combustion and the molar heat of combustion of pure paraffin is
between 0 and 350 kJ/mol).
8.
Methyl alcohol, CH3OH(1), is generally used as a fuel to heat the oil in fondue
sets. What mass of methyl alcohol is needed to heat 1.10 kg of fondue oil from
22.0°C to 98.0°C?
Specific heat capacity of fondue oil is: 8.9 kJ/[kg°C]
Equation for the combustion of methyl alcohol is:
CH3OH(1) + 3/2 O2(g)  CO2(g) + 2 H2O(g) + 639 kJ
9.
Styrofoam
Thermometer
Water
During a calorimetry experiment, 4.01 g of LiCl(s) are dissolved in 100. mL of
water. The temperature of the water increased from 22.0°C to 28.0°C. What
quantity of heat energy would be produced if 2.0 moles of LiCl were dissolved?
45
II. Thermochemistry
10.
A blacksmith heats a horseshoe to a very high temperature; the horseshoe has a
mass of 550 g.
Then he plunges it into an insulated bucket containing 8.50 liters of water at
22.0°C.
He notes that the maximum water temperature reached was 26.9°C.
Given that the specific heat of the iron is 0.45 kJ/[kg °C], determine the
temperature to which the horseshoe must have been heated.
Extra Calorimetry (solutions also on web site)
1.
Phileas Fogg, the character who went around the world in 80 days, was very
fussy about his bathwater temperature. It had to be exactly 38.0 o C. You are
his butler, and one morning while checking his bath temperature, you notice
that it’s 42.0oC. You plan to cool the 100.0 kg of water to the desired
temperature by adding an aluminum-duckie originally at freezer temperature
(-24.0oC). Of what mass should the Al-duckie be? [Specific heat of Al = 0.900
J/(goC); density of water =1 .00 g/ml]. Assume that no heat is lost to the air.
2.
A certain material’s (environment) temperature increases by 1.00oC for every
1560 J that it gains. A 0.1964 g sample of quinone (molar mass = 108.1
g/mole) was burnt, and the surrounding material’s temperature increased from
20.3 oC to 23.5 oC. Find the molar heat of combustion for quinone.
3.
A 1.55 g of CH4O sample is burnt in a calorimeter. If the molar heat of
combustion of CH4O is -725 kJ/mole, and assuming that the 2.0 L of water
absorbed all of the heat of combustion, what temperature change did the water
experience?
4.
0.20 moles of HX were neutralized by NaOH. The concentrations of the base
and acid were equal. If the temperature of the water in the calorimeter
increased from 19.9 to 24.6 C, what was the original concentration of HX?
Molar heat of neutralization = -80 kJ/mole of HX
5.
In real calorimeters, most of the heat released by the bomb is absorbed by
water, but a certain amount is also absorbed by the metal and insulation
surrounding the water tank. A certain calorimeter absorbs 24 J/oC. If 50.0 g of
52.7oC water is mixed with the calorimeter’s original 50.0 g of 22.3oC water,
what will be the final temperature of the mixture?
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