Download Chapter 7 – Chemical Reactions and Energy Flow

Chapter 7 – Chemical Reactions
and Energy Flow
Energy is the currency of the Universe. Chemical changes
are usually accompanied by a redistribution of energy
and not just matter. And whether or not a reaction will
proceed depends upon the flow of energy.
The study of the energy flow is called “thermodynamics” –
literally, “heat movement” – but initially, we will be
considering a subset called “thermochemistry”. The two
topics are related and often intermixed.
Chapter 7
1
Some reactions release heat when they occur; others do
not. Indeed, some reactions absorb heat. How do we
account for these observations?
The model that we have developed is that energy is
stored in chemical compounds and if the energy stored in
the reactants is greater than the energy in the products,
then some will be left over and released. (And vice versa).
We call processes that release energy “exothermic”
(“exo” means out or outside, so “exothermic” means
“heat out”). Processes that absorb energy are then
“endothermic” (“endo” means in or inside).
Chapter 7
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Q7.1)
The combustion of a candle is:
a) exothermic
Q7.2)
c) neither
The melting of an ice cube is:
a) exothermic
Q7.3)
b) endothermic
b) endothermic
c) neither
The Kreb Cycle is:
a) exothermic
b) endothermic
Chapter 7
c) neither
d) who knows?
3
During a chemical reaction, a shift from reactants to
products will result in the release or absorption of energy.
In an exothermic process, the initial state has more energy
than the final state and the excess energy is released as
heat. In an endothermic process, the final state has more
energy than the initial state and heat is absorbed.
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Energy is the capacity to do work – to apply a force over a
distance.
Kinetic energy is the energy tied up with motion – the faster
a molecule is moving, the more kinetic energy it has.
Potential energy is the energy of place – for molecules, this
means the relative amount of energy invested in the bonds
within a molecule and the attractive/repulsive forces
between atoms, ions, and molecules.
Heat is manifested in changes in both kinetic and potential
energy within a system and its surroundings.
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Heating a container of water results in:
1) A change in kinetic energy. The
individual molecules of water accelerate
and move much rapidly within the solution.
Some have sufficient kinetic energy that
they are capable of overcoming their
intermolecular attractions and leave the
solution.
2) A change in potential energy. The molecules, on
average, move further apart (which is why the density
decreases with increasing temperature) and work is
required to move the molecules against the forces of
attraction.
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This conversion of energy from one
form to another occurs freely in all
sorts of different systems. For
example, chemical energy can be
converted to electrical energy which
manifests in the form of both light and
heat (electromagnetic radiation).
The “law of conservation of energy” or the “first law of
thermodynamics” tells us that although energy can be
converted from one form to another, the total energy is a
constant.
To put it another way “you can never win” – you can
never get more energy out of a system than is present in
the system to begin with.
Chapter 7
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Energy units
Energy is most often measured – in scientific terms – in
“joules” which is the energy required to lift one kilogram
one metre. A joule is a “kg•m2/s2” or “kg m2 s-2”.
An older unit of heat energy was the “calorie” which is
defined as the amount of energy required in order to raise
the temperature of 1 gram of water by 1°C. The conversion
factor is 1 cal = 4.184 J (recall the specific heat capacity of
water).
This should not be confused with the “food calorie” (Cal.)
which is actually a “kilocalorie” or the amount of food
needed to raise one kilogram of water by 1°C. Just
keeping our bodies warm requires about 2000 Cal per day.
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A second consideration that is important in
thermochemistry is “what exactly are we measuring?”
Consider that the “law of conservation of energy” tells us
that energy can be neither created nor destroyed – that
the amount of energy is constant. If that is the case, then
how can we say that a reaction is “exothermic”? That it
“releases energy”. This would imply that the reaction
creates energy that is released.
The answer is that in thermochemistry, we are talking about
a system and when considering whether a reaction is
“exothermic” or “endothermic”, we are considering
whether energy is leaving or entering the system. That total
energy is a constant.
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We define the system as the
thing that we are studying – a
chemical reaction, the beaker
in which it occurs, the
instruments monitoring it or
whatever. We then define the
surroundings as everything else
in the Universe.
System + Surroundings = Universe
energy of a system
+ energy of surroundings = constant
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An open system is one in which there is the free exchange
of both heat and matter between the system and
surroundings. Think of a pot of boiling water on a stove.
Heat is entering from the stove’s element and matter is
leaving in the form of steam (water vapour).
A closed system is one in which heat is freely exchanged
but matter is not. Think of a pressure cooker filled with
water. The lid ensures that even when boiling, none of the
water is released. (Most of the time, this is what we are
talking about….)
An isolated system is one in which neither heat nor matter
is exchanged with the surroundings. Doesn’t actually exist
– except in the sense that the entire Universe is an isolated
system. However, theoretically this is a useful concept.
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Q7.4)
a) open
Q7.5)
a) open
Q7.6)
a) open
What type of system is a Duracell non-rechargeable battery?
b) closed
c) isolated
d) not a system
What type of system is a Energizer rechargeable battery?
b) closed
c) isolated
d) not a system
What type of system is the Krebs cycle?
b) closed
c) isolated
Chapter 7
d) not a system e) who knows?
12
So, heat is a measure of the flow of energy between the
system and its surroundings. An object or solution or
chemical compound does not contain heat. Rather, there
is thermal energy within a system that can manifest as
heat. Wood, for example, has less thermal energy than
coal on a mass bases. But in either case, reacting the fuel
with oxygen results in a lot of heat to the surroundings.
Temperature refers to how hot something is – to the
average kinetic energy of the molecules or atoms or ions
within a system. A system at a high temperature can
transfer energy in the form of heat to surroundings that are
at a lower temperature.
Temperature is an intrinsic property – it doesn’t depend on
how much substance is present. Heat is not intrinsic.
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Heat transfer occurs between any two objects when they
come in contact until they reach “thermal equilibrium”.
Heat is said “to flow” from the object with the higher
temperature to the object with the lower temperature.
This is an important distinction, because it is different from
the way we view this in our daily lives. We often talk about
“the cold creeping in” when in reality, it is the heat flowing
out. That is, holding an ice cube results in the loss of heat
from your hand to the ice cube not the cold from the ice
cube flowing into your hand – although that is what
appears to be happening.
Thermal equilibrium is achieved when both objects have
the same thermal energy – and the same temperature.
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The energy content of a system is the sum of all of the
kinetic and potential energy in the molecules, atoms, or
ions within a system. This is referred to as the “internal
energy (U)”. The internal energy depends on the amount
of matter (it is an “extrinsic property”) as each ion, atom, or
molecule in a system contributes to the total.
Internal energy is also a “state function” – a function that
depends only on the state of a system and not on its
history.
We can’t measure the internal energy of system – but we
can measure the change in the internal energy:
ΔUsystem = Ufinal -
Chapter 7
Uinitial
15
Heat and work are both forms of energy transfer between
a system and its surroundings. The transfer of heat results in
a change in temperature. Work must be accompanied by
a change in volume against a constant pressure. (Work
requires the expenditure of energy over a distance.)
The change in internal energy is the sum of the energy
transferred as heat and work:
Δusystem = q
+
w
where “q” is heat and “w” is work.
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Note that if we are talking about the amount of heat (i.e.
102 J) then we don’t need to specify a sign. But when we
are accounting for internal energy, we define q and w as
positive if they increase the internal energy of a system.
Most of the time, the direction or sign is apparent from the
language used.
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In a system at constant pressure – a beaker open to the
atmosphere – we define a slightly different quantity for our
discussion of energy.
Enthalpy (H) is the amount of heat transferred between a
system and its surroundings during a process that occurs at
constant pressure (qp), if no work other than that due to
expansion of the system occurs.
Enthalpy is a state function:
ΔH
=
Hfinal -
Hinitial
For exothermic processes, at constant pressure, ΔH < 0. For
an endothermic processes, at constant pressure, ΔH > 0.
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“Fourth Law of Thermodynamics: If the probability
of success is not almost one, then it is damn near
zero.”
- David R. Ellis
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Note that if we are talking about the amount of heat (i.e.
102 J) then we don’t need to specify a sign. But when we
are accounting for internal energy, we define q and w as
positive if they increase the internal energy of a system.
Most of the time, the direction or sign is apparent from the
language used.
Chapter 7
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In a system at constant pressure – a beaker open to the
atmosphere – we define a slightly different quantity for our
discussion of energy.
Enthalpy (H) is the amount of heat transferred between a
system and its surroundings during a process that occurs at
constant pressure (qp), if no work other than that due to
expansion of the system occurs.
Enthalpy is a state function:
ΔH
=
Hfinal -
Hinitial
For exothermic processes, at constant pressure, ΔH < 0. For
an endothermic processes, at constant pressure, ΔH > 0.
Chapter 7
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Chapter 7
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Changes in state require energy. Boiling water requires the
input of heat but what is less obvious, perhaps, is that the
condensation of steam releases an equivalent amount of
heat.
The energy involved is labelled for the conversion
involved. The “molar enthalpy change of fusion” (ΔHfus) is
the energy involved in melting/freezing – the conversion
from a liquid to a solid and vice versa. The “molar
enthalpy change of vapourization” (ΔHvap) is the energy
involved in boiling/condensation. There is also a “molar
enthalpy change of sublimation” (ΔHsub) which is the
energy involved in sublimation/deposition.
Chapter 7
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Chapter 7
24
For water:
ΔHfus = 6.00 kJ/mol or 333 J/g; ΔHvap = 40.65 kJ/mol or 2256 J/g
specific heat of ice: 2.06 J/K·g water: 4.184 J/K·g steam: 1.92 J/K·g
How much heat is required to take a 225 g block of ice from -4°C to 37°C?
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We can talk about the change in enthalpy for a reaction.
It is defined as the difference in the sum of the enthalpy of
the products minus the sum of the enthalpy of the
reactants:
ΔHrxn =
Σ Hproducts - Σ Hreactants
However, we can’t actually use this equation because
we can’t measure the values of the absolute enthalpies
of the products or reactants.
But we can measure the change in enthalpy of a
reaction using a “bomb calorimeter”. We can also
calculate Hrxn using the “molar enthalpies of formation”.
Chapter 7
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Enthalpy diagrams for
exothermic and endothermic
reactions show the qualitative
relationship between the
reactants and products in a
chemical reaction. Such
diagrams can help visualize
energy changes.
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27
But, ideally, we would like to work with quantitative
relationships – how much energy can be get from a
reaction?
For example:
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g) ΔHf = -890.3 kJ/mol
In this case, the enthalpy of the reaction is the amount
released per mole of methane used. (Book mentions
“packets” – which is one way to think of it.)
The magnitude of the enthalpy change depends on the
amount. If we were to react 2 moles of methane, we
would get 2 mol x -890.3 kJ/mol or -1780.6 kJ. For 0.5 moles
of methane, we only get -445.15 kJ.
Chapter 7
28
The sign depends on the direction of the reaction. Turn the
reaction around:
CO2(g) + 2H2O(g)
CH4(g) + 2O2(g) ΔHf = +890.3 kJ/mol
Which makes sense from our qualitative
picture. If going one direction gives up
energy, going the other direction
requires the input of energy.
Note that part of what it means to be a state function is
that the reaction will generate the same amount of
energy provided that methane is at 1 bar and 25ºC,
regardless of the history of the methane or where the
reaction takes place.
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Calorimetry is a fairly simple
method for measuring the
enthalpy of a reaction. In
principle, (and in first year
labs) a calorimeter is simply a
device (i.e. coffee cups and a
thermometer) that isolates a
reaction in a way that allows
for a reproduce-able
measurement of the amount
of heat generated or
consumed during a reaction
The amount of heat is obtained from the change in the
temperature of the water in the calorimeter.
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We know the specific heat of water is 4.184 J/K•g and this
leads to the expression:
q = c x m x ΔT
where “m” is the mass of water and “c” is the specific
heat. For example, 50 grams of water, with a temperature
change of 3.2ºC, requires:
4.184 J/K•g x 50 g x 3.2 K = 669.4 J
The accuracy of the calorimeter depends upon the
accuracy of measuring the temperature and, in practical
terms, coming up with a specific heat capacity for the
calorimeter itself to account for heat losses. This is
accomplished by using a reaction of known values to
calibrate the device.
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31
Instruments used in research laboratories (and upper level
undergraduate laboratories) are a little more
sophisticated. But they operate on the same principles
and need to be calibrated before using.
Chapter 7
32
Example:
A bomb calorimeter is calibrated by combusting 1.06 g of
benzoic acid in excess oxygen resulting in a 4.63ºC
temperature rise. A 1.83 g sample of an unknown is combusted
resulting in a 3.86ºC rise. If the heat of combustion for benzoic
acid is 26.43 kJ/g, what is the heat of combustion of the
unknown?
First, calculate “c” for the calorimeter:
q = (26.43 kJ/g)(1.06 g) = 28.0158 kJ
produces a 4.63ºC change which means that a temperature
change of 3.86ºC is produced by:
q = (3.86ºC/ 4.63ºC) x 28.0158 kJ = 23.3566 kJ
which was produced by 1.86 g, so the heat of combustion is:
ΔHcombustion = 23.3566 kJ/1.83 g = 12.76 kJ/g
Chapter 7
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If we are going to compare compounds or reactions with
one another, we need to have a “standard state” – since
enthalpy depends on pressure, concentration,
temperature, etcetera.
Standard State:
- For pure substances, the standard state is the most
stable form and state of the substance at 1 bar at the
temperature of interest (for most compounds, 25ºC).
- For any gas, its state at a pressure of 1 bar.
- For any aqueous species, its standard state when the
concentration is exactly 1 M at a pressure of 1 bar (and
usually 25ºC).
For a reaction with reactants and products in their
respective standard states, we have the standard enthalpy
change of reaction.
Chapter 7
34
We can utilize the standard enthalpy of reaction to work
out the energy for any reaction – or any sequence of
reactions.
Hess’ Law: If a reaction can be written as the sum of two or
more steps, the enthalpy change of the overall reaction is
the sum of the enthalpy changes of the reactions of the
individual steps.
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Hess’s Law is a necessary consequence of the law of
conservation of energy. The total energy of the steps must
be the total energy of the overall reaction.
Note that you must take stoichiometry into account in
calculating the steps for a Hess’s Law calculation.
Chapter 7
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“Any intelligent fool can make things bigger,
more complex, and more violent. It takes a touch
of genius -- and a lot of courage -- to move in the
opposite direction.”
- Albert Einstein
Chapter 7
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The enthalpy of a reaction is equated with the heat that
is given off from a reaction. It can be measured with
instrumentation (a calorimeter). It can be assessed for
reactions using Hess’s Law.
It can also be determined using the “standard molar
enthalpy change of formation” (ΔHf) which is:
“the enthalpy change accompanying the generation of
a compound in its standard state from its constituent
elements in their standard state and their stoichiometric
proportions. The temperature must be specified.”
This works because all of the compounds in a reaction
are then measured against a common point – their
elements.
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For example, consider the combustion of sugar – sucrose.
C12H22O11(s) + 12O2(g)
12CO2(g) + 11H2O(g)
ΔHrxn = -5645.1 kJ/mol
The ΔHf of sucrose, carbon dioxide, and water are all
required. Note that oxygen is the elemental form of
oxygen and its standard state at 25ºC with ΔHf = 0 kJ/mol.
To calculate the other three values, we measure the
reactions between the chemical species and their
elements:
12C(s) + 12O2(g)
12CO2(g)
11H2(g) + 5.5O2(g)
11H2O(g)
C12H22O11(s)
12C(s) + 11H2(g) + 5.5O2(g)
Chapter 7
39
Individually, these reactions are:
C(s) + O2(g)
CO2(g) ΔHf = -393.51 kJ/mol
H2(g) + O2(g)
H2O(g) ΔHf = -285.83 kJ/mol
C12H22O11(s)
12C(s) + 11H2(g) + 5.5O2(g)
ΔHf = +2221.2 kJ/mol
But we need to take the stoichiometric factors into
account, so the enthalpy for the overall reaction is then:
ΔHrxn = 12 x -393.3 kJ/mol + 11 x -285.83 kJ/mol
+ 1 x 2221.2 kJ/mol + 12 x O kJ/mol
= -5645.1 kJ/mol
Chapter 7
40
The standard enthalpy
values for many
substances are
tabulated. Appendix C
has a fairly extensive list.
Chapter 7
41
The standard enthalpy of a reaction at any specified
temperature can be defined by:
ΔHºrxn
=
Σ ni ΔHºf (products) - Σ ni ΔHºf (reactants)
Note that this is what we did with sucrose except here we
are subtracting the total for all of the reactants and we
use the signs, as written, in the tables provided.
(subtracting a negative number is the same as adding a positive….)
Also, we don’t actually need to write out all of the
equations….
Chapter 7
42
Suppose we want to calculate the enthalpy change for
the decomposition of calcium carbonate to calcium
oxide and carbon dioxide at 25ºC. What is the value?
CaCO3(s)
ΔHºrxn
=
CaO(s) +
CO2(g)
Σ ni ΔHºf (products) - Σ ni ΔHºf (reactants)
= [1x ΔHºf(CaO) + 1x ΔHºf (CO2)] – [1x ΔHºf (CaCO3)]
= [(-635.1 kJ/mol) + (-393.5 kJ/mol)] – [(-1207.6 kJ/mol)]
= + 179.0 kJ/mol
Note that this implies that it takes energy to separate
carbon dioxide from limestone – which is why lime won’t
work to sequester CO2.
Chapter 7
43
Despite extensive tables containing the standard
enthalpy of formation for literally hundreds of chemical
compounds, they are not sufficient. Many reactions are
outside of the scope of the tables.
However, we can approximate the energy by using the
average bond energies for all of the chemical bonds
involved. That is, if we know the amount of energy
involved in each bond, we can work out an approximate
value for each chemical species and work out the
enthalpy of reaction from the difference.
The bond energy (D) is “the enthalpy change for breaking
a particular bond in the molecules, assuming the
reactants and products are in the gaseous state.”
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44
The energy difference is the energy of the bonds broken
minus the energy of the bonds made:
ΔHºf =
Σ D(bonds broken)
-
Σ D(bonds made)
This is simple a case of drawing an appropriate structural
formula and then working through the bonds in the
molecule:
H3C-CH3
H3C
6 x H-C
6x [ H
CH3
ΔHºf = D = +346 kJ/mol
C ] ΔHºf = D = 6x +413 kJ/mol
D = 2824 kJ/mol
The bond energies are tabulated.
Chapter 7
45
Chapter 7
46
Calculating the enthalpy of combustion of ethane gives:
CH3CH3(g) + 3.5O2(g)
For oxygen:
For carbon dioxide:
For water:
2CO2(g) + 3H2O(g)
D = 498 kJ/mol
D = 2x 783 kJ/mol = 1566 kJ/mol
D = 2x 463 kJ/mol = 926 kJ/mol
Which gives:
ΔHºf = (2824 kJ/mol + 3.5x498 kJ/mol)
– (2x1566 kJ/mol + 3x926 kJ/mol)
= -1343 kJ/mol
________________________________________________________________
Note using the enthalpies from Table 7.2 gives:
ΔHºf = (2x-393.51kJ/mol + 3x-285.83kJ/mol)
-(1x-83.85kJ/mol + 3.5x0.0kJ/mol) = -1428.9 kJ/mol
Chapter 7
47
All sorts of reactions can yield useable energy but
combustion is the most obvious. Burning propane and
ethane light up our barbecues; burning wood provides the
warm glow of a campfire; burning gasoline drives our cars;
and burning food warms our bodies while providing the
energy to live.
The many different chemical reactions that food undergo
in our bodies to provide energy and to provide the building
blocks for many of our biomolecules fall into the area of
biochemistry called “metabolism”.
Primary amongst these reactions is the combustion of
glucose (a carbohydrate):
C6H12O6(s) + 6O2(g)
6CO2(g) + 6H2O(l) ΔHºrxn = -2803 kJ/mol
Chapter 7
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Of course, when we refer to this as “combustion”, we really
are saying “oxidation” – the carbon atoms in the glucose
are oxidized to carbon dioxide. This is the underlying
principle behind any combustion reaction. Further, if we
were to use this energy directly, we would need to
continuously stoke the fires.
Instead, our bodies have evolved a mechanism that is
common across all of life. We convert the energy of
combustion from one mole of glucose to the formation of
30-32 ATP (adenosine-5’-triphosphate) molecules via
aerobic respiration.
ATP then undergoes a hydrolysis reaction (literally, “water
breaking”) to produce ADP (adenosine-5’-diphosphate).
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49
ATP(aq) + H2O(l)
ADP(aq) + HPO42-(aq) ΔHºrxn ≈ -24 kJ/mol
Chapter 7
50
Note that ATP doesn’t have “high
energy bonds”. Rather, the total
energy of the reactants is greater
than that of the products. And that
energy must be loaded back into
the system by the combustion of
glucose.
Also, that glucose is ultimately
formed from the reaction of carbon
dioxide and water in photosynthesis
– the reverse of the combustion
reaction of metabolism.
Chapter 7
51
“A good working definition of infinity is
waiting for something to be perfect.”
- Tony Donovan, EMS
Chapter 7
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