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1218 Journal of Chemical Education • Vol. 85 No. 9 September 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Classroom
In recent years, there has been a popular interest in oxygen
as it has become the name of a Broadway play, a new health
magazine, a television network, a computer project at MIT,
and at least one new book. In spite of this, and because of its
availability, its importance as an energy source is commonly
overlooked. It is more common to speak of fossil fuels, hydro-
gen, or carbohydrates as the major energy sources in chemistry
and biology but these compounds are some of the most stable
organic molecules found in nature. ese compounds are made
of strong bonds and have little tendency to combine with any
other molecules. Indeed, hydrocarbons were known as parans
because of their low anity for reaction. However, oxygen is a
diradical held together with a bond energy of just 496 kJ mol‒1
(1). By contrast, the oxygen atom forms two strong bonds in
each of the combustion products; carbon dioxide (799 kJ mol‒1)
and water (926 kJ mol‒1). As a result, every mole of oxygen that
reacts releases, on average, 460 kJ.1 It is the uniquely weak bond-
ing in the oxygen molecule that makes combustion one of the
most exothermic reactions in chemistry. By contrast, the bond
strengths of the organic molecules are very similar to the bond
strengths of the combustion products. us, in a chemical sense,
the oxygen molecule is the energy source and the other “fuels”
are merely vehicles to allow each oxygen atom to form strong
bonds in the combustion products. It is the relative scarcity of
compounds to react with the plentiful oxygen that leads to an
underestimation of the importance of oxygen.
A striking result of this fact is that the heat of combustion
of any organic molecule can be calculated approximately by
merely balancing the reaction and determining how many moles
of oxygen are reacted when the organic molecule is burned. For
example, the oxidation of methane uses two moles of oxygen per
mole of methane releasing approximately 2 × 460 = 920 kJ mol‒1
of methane. e experimental heat of combustion of methane is
890 kJ mol‒1 (2). For a mole of glucose, which combines with
6 moles of oxygen to form CO2 and water, the calculated value
is 2760 kJ whereas the experimental value is 2801 kJ mol‒1 of
glucose (3). erefore the approximate heat of combustion (or
partial oxidation) of any organic molecule may be calculated
with reasonable accuracy simply by balancing the equation. Of
course, one can also calculate the caloric value (Calories per
gram) of sucrose or other common foods from this approach.
e breadth of these calculations can be seen in Table 1. While
these values are not exact, the conceptual value of these calcula-
tions should be clear.
e role of oxygen as the energy source is even more obvi-
ous in biological systems. e majority of ATP bonds generated
in glucose metabolism derive from the oxidation of the hydro-
gen atoms that are removed from the organic substrates. e
aerobic metabolism of a glucose molecule consumes six oxygen
molecules and can generate 38 ATP bonds. is should raise the
question about the origin of the energy produced in anaerobic
processes. With no available oxygen, the anaerobic metabolism
of glucose can generate only two molecules of ATP per glucose
molecule. e small quantity of energy required for these two
ATP bonds comes from rearranging the oxygen atoms of the
carbohydrate molecule to generate products with improved
resonance structures. Note that no oxygen is consumed in these
two balanced reactions (Scheme I). e intuition gained from
this approach can be very valuable for students.
Many textbooks point out that our fossil fuels and carbo-
hydrates derive their chemical energy from the sun. It should be
noted that methane is abundant on many planets and moons
far from the sun but that free oxygen is found on none of them.
ere are plenty of oxygen atoms out there but they are all bound
up with carbon, hydrogen, metals, silicon, and so forth. While
the earthly “fuels” may result from the sun’s energy, it is really
the oxygen molecules generated during photosynthesis that are
trapping the energy of the sun and facilitating life on earth.
Given the great exothermicity of these oxygen reactions, it
is surprising that so many organic compounds can survive their
Appreciating Oxygen
Hilton M. Weiss
Department of Chemistry, Bard College, Annandale-on-Hudson, NY 12504; weiss@bard.edu
CH
OH
OH
C
O
3
HC
3
HC2
2 2
3
HC CH2OH
C
O
O
C
O
O
C
H O
CHOH
CHOH
CHOH
CHOH
CH OH
2
muscle
yeast
CH
OH
C
O
OH
Scheme I. Balanced equations of the anaerobic metabolism of
glucose.
Table 1. Enthalpy Calculated from the Amount of Oxygen Used
Compound Moles
of O2
ΔHo/(kJ mol–1) Error (%)
Calc Exp
Methane 2.0 –920 –890 3.4
Octane 12.5 –5750 –5452 5.5
Methanol 1.5 –690 –726 5.0
Ethanol 3.0 –1380 –1367 1.0
Benzoic Acid 7.5 –3450 –3227 6.9
Sucrose 12.0 –5520 –5644 2.2
Thiophene 6.0 –2760 –2805 1.6
No t e : The experimental data is from ref 2.
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 9 September 2008 • Journal of Chemical Education 1219
In the Classroom
constant exposure to oxygen. Why don’t we experience “spon-
taneous human combustion”? Put another way, if oxygen is so
thermodynamically reactive, why is it so kinetically unreactive?
is is the second great property of oxygen. Because its two
unpaired electrons have the same spin, the rst step of its reac-
tion with any diamagnetic molecule (having paired spins) must
generate two radical products with the same spin. ese initial
radical products are usually high energy reactive intermediates
and their formation will be slow. Once the reaction is initiated
however, subsequent reactions can easily continue the chain of
further reactions that lead to carbon dioxide and water. Anti-
oxidants are compounds that can stop these chains by forming
stable radical products.
us, oxygen is the unique molecule that stores the energy
coming from the sun and provides the energy for life on earth.
Note
1. e value of 460 kJ mol‒1 is an approximate value averaged
from calculations using experimental data of numerous organic
compounds. For example, using the data in Table 1, the experimental
∆H° for octane is –5452 kJ mol‒1. Balancing the equation for the
combustion of octane requires 12.5 moles of oxygen. erefore this
reaction evolves –5452/12.5 = –436 kJ per mole of oxygen reacted.
Methanol needs 1.5 moles of oxygen and evolves 720 kJ of heat, which
gives –720/1.5 = –484 kJ per mole of oxygen. e average value is
(436 + 484)/2 = 460 kJ per mol of oxygen. Other examples give values
closer to 460 kJ but an average value is in this neighborhood.
Literature Cited
1. Bond energies are taken from Atkins, P.; Jones, L. Che mistry:
Molecules, Matter, and Change, 3rd ed.; W. H. Freeman: New
York, 1996.
2. Noggle, J. Physical Chemistry, 2nd ed.; Scott, Foresman and Co.:
Boston, 1989; p 279.
3. Chang, R. Chemistry, 7th ed.; McGraw-Hill: New York, 2002; p
217.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2008/Sep/abs1218.html
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