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Energy in Bonds
Name: Unknown
Status: educator
Age: 40s
Location: N/A
Country: N/A
Date: 1999-2001
Question:
How is it that energy can be stored in chemical bonds if
creating bonds releases energy and breaking bonds
requires energy? These two facts seem to contradict each other.
Replies:
I would use a simple example of the Hindenburg, a
dirigible which was filled with hydrogen and was
therefore 'lighter than air'. The craft was lost in a
violent explosion when it hit electric cables which
provided the spark to get things going. The explosion
consumed the hydrogen and surrounding oxygen , in the
process producing water, with a great NET RELEASE of
energy. Note that even though the spark was needed in
the beginning to begin the process, there was far more
energy released in the formation of the bonds of the
resultant water molecules.
You might want to consult a good chemistry text for an
explanation of exothermic and endothermic reactions.
Not all chemical reactions produce a net release of
energy; some require a greater input of energy than is
yielded in the end.
Thanks for using NEWTON!
Ric Rupnik
There is no contradiction. Let us walk through the processes. If chemical
species 'A' and 'B' react to form molecule 'A-B', let us say a certain amount
of energy, Eex [standing for an ex-othermic reaction], is given off. If we
carry out the opposite process 'A-B' react to break the bond giving the
species 'A' and 'B' a certain amount of energy is required to break that
bond, call it, Een [standing for
an en-dothermic reaction]. Eex is negative and Een is positive by convention
but they have the same magnitude. There is no contradiction in that
description of the formation and breaking of a bond.
In biology, energy is supplied by the reaction of complex species. That
energy comes from light by photosynthesis, or from other substances by the
digestion of foods. The energy supplied by these processes is "stored" in
various chemical substances until the plant or animal "calls for them" to be
used in some biological function. These reservoirs of energy then react,
giving off energy, if that is the function the organism is "demanding", to
form other chemical species called
"metabolites". These latter species then are eliminated from the organism,
be it oxygen in the case of plants, carbon dioxide and other species, in the
case of animals.
This "bio-energy, life force" does not appear by magic. It does not
contradict the conservation of energy. The energy comes from a source
outside the organism and is just stored in a chemical form that can be
called upon as needed.
Vince Calder
There is a natural tendency for all "systems" to end at the lowest free
energy state (G...Gibbs free energy...actually a restatement of the second
thermodynamic "law" all systems tend toward an equilibrium. This is just
like a ball falling to the ground.
On the other hand all spontaneous reactions will have a negative (G.
Consider the hydrogen bond...in its formation there is a release of about 25
kJ/mole and at the same time there is an almost equivalent loss in entropy.
So much for "chemistry"...Now in Biological systems they are fighting a war
against entropy and loss of free energy. The biological system must expend
constant energy or be reduced to a state of equilibrium and disorder
equaling death. Biological systems absorb either energy directly from the
sun or simple chemical or from complex chemicals to generate temporary
stores of usable energy for driving reactions with a positive (G. The
statement that a bond is the "lowest potential" is not true because some
bonds like ATP~P~P are unstable high-energy bonds that are easily broken.
In other cases, as in the saturation of carbon chains to form lipids (fats
and oils), the bond energy can be utilized and ripped off in a orderly
sequence to be used to drive reactions with a +(G. I hope this didn't make
it more confusing.
pf
Peter Faletra Ph.D.
Office of Science
Department of Energy
Bond breaking always requires energy input. Otherwise everything would fall
apart all by itself. Bond making always releases energy. It is the energy
interplay between that released and that required which drives chemical processes.
Getting specific: There is simply no such thing as a high energy bond in ATP.
Severance of the terminal phosphate group requires energy. It is the
production of the inorganic dihydrogen phosphate ion -- a bond making process
-- that accounts for the apparent energy release of 7.3 Kcal/mol when the
terminal phosphate of ATP is hydrolyzed.
That the remaining phosphate bonds in ADP and AMP respectively "release" less
energy is explained by the fact that the bonds to be hydrolyzed in ADP and AMP
are stronger than the one broken in ATP. More energy is required to break them
and correspondingly less energy is derived from formation of the product
species.
Consider the familiar (exoenergetic) neutralization of NaOH with HCl: The heat
released comes not from breaking apart the Na+ and OH- or from splitting the
H+ from the Cl- -- the energy is derived from the formation of water
then the H+ and OH- combine a bond making process. Same is true when .
Energy rich substances are those that required a lot of energy to make them.
These are substances that release a lot of energy when they are converted to
more stable (lower energy) stuff. Consider nitroglycerin. It is made of
nergy
poor elements, C, H, O. and N. Driven through a sequence of energy
"increasing" reactions, these elements emerge as a very unstable material
which can be easily converted to vastly more stable (energy poor) stuff --
CO2, H2O, and NOx.
Sincerely,
ProfHoff
All of what you have said is true, i.e., "creating bonds releases
energy and breaking bonds requires energy". We are, of course, talking
about the formation of a thermodynamically stable molecule, and the bonds
which are formed within this stable molecule are, indeed, in a low energy
state. This is not to say, however, that we could not take the same atoms
that make up this molecule, and put them together into one or more other
molecules that are even more stable. For example, TNT is a stable molecule,
but could have formed several other molecules from the carbon, hydrogen,
oxygen, and nitrogen atoms that we used to manufacture this TNT molecule.
In fact, nature, if it had its choice, would have made other molecules from
these atoms, exactly because these other molecules are more stable. In
other words, if we add up the total energy required to put together these
more stable molecules, we would see that it takes less energy to do this
than it does to make the TNT molecule. Therefore, if we give nature a
chance to rearrange the atoms within the TNT molecule (by supplying some
energy), these atoms will spontaneously reassemble themselves into the
group of several, more stable molecules. In fact, most explosives which you
are familiar with rely on the release of energy when oxygen-nitrogen bonds
within explosives molecules are broken. The resulting molecules which are
formed from the constituents of the original explosives molecules, things
like water, carbon dioxide, and nitrogen gas, do require energy to make,
but this energy cost represents only a small fraction of the energy which
is required to make the TNT molecule. The excess energy, liberated over a
short period of time, is what gives explosives their great power.
I hasten to point out that the excess energy that is recovered by
allowing the atoms within the TNT molecule to rearrange themselves into the
group of more stable atoms will not be 100% of the theoretical amount. All
chemical reactions are lossy, in that some of the energy is converted into
forms of energy which cannot be used to do useful work. To quote an author
(whose name I have sadly forgotten), the Laws of Thermodynamics can be
paraphrased as
1st Law : "You can't win the game"
2nd Law : "You can't break even"
3rd Law : "You can't ever leave the game"
The bottom line, is one of relative stability. Remember, the
word "stable" is a relative term.
Jim Rubin
You are quite correct; breaking a chemical bond requires energy.
Chemically, when we speak of "bond breaking," we mean homolytic cleavage, in
which the two fragments each contain half of the electrons that were in the
broken bond. The products of this process always have higher energy than
the starting material with an intact bond. However, this sort of process
isn't really relevant for biological systems.
When biology texts state that certain chemical units, usually phosphate
esters or di- and tri-phosphates, contain "high-energy bonds," that is
shorthand for saying that they are "high-energy" when compared to some other
kind of bond, not compared to broken bonds. So the statement isn't
literally true, but it actually is useful.
As you know, living systems use certain functional groups, particularly
phosphates, to transfer energy and to enable molecular synthesis. Let's
take as an example the main energy currency of biochemistry: adenosine
triphosphate, ATP. Many biochemical processes are driven by coupling to the
hydrolysis of ATP, which is often conveniently written as
ATP --> ADP + Pi + energy.
That is, ATP breaks into ADP and an inprganic phosphate. Sometimes the
"high-energy phosphate bond" of ATP is used to explain how this reaction
produces energy. It's actually a little more complicated than that.
ATP contains oxygen links between phosphorus atoms, P-O-P. Similarly, other
"high-energy" species often found in biological cycles, such as
glucose-6-phosphate, contain oxygen links between carbon and phosphorus,
C-O-P. These "pyrophosphates" and "phosphate esters" can be transformed,
with the aid of enzymes, into other substances with lower energy. The
difference between the energy of the starting materials and the energy of
the products is the energy released in the transformation.
In the case of pyrophosphates and phosphate esters, another important
reactant is water. A water molecule can be inserted between the two
phosphorus atoms of a pyrophosphate, decomposing the structure:
P-O-P + H-O-H --> P-O-H + H-O-P + energy.
A phosphate ester can do something similar:
C-O-P + H-O-H --> C-O-H + H-O-P + energy.
(In both of these cases, there are other atoms connected to the carbon and
phosphorus atoms. I have omitted them here to concentrate on the action and
because it's hard to draw three-dimensional molecules in a one-dimensional
line of text.) These reactions are often written in biology texts in
shorter form as
P-P --> P + Pi and
C-O-P --> C-OH + Pi.
But you need to remember that the end products don't have broken bonds.
Instead, some of the bonds are replaced by other bonds, resulting in lower
energy products.
So, whenever you see an illustration of energy being released by cleaving
"high-energy bonds," just remember that what is actually happening is that
P-O-P and C-O-P linkages are being replaced by P-O-H and C-O-H linkages.
Does that mean they high-energy bonds are being replaced by low-energy
bonds? In a way, yes. That's not a bad way to think about it. They're
just leaving out the water molecules and the full structure of the inorganic
phosphate because it's shorter to write and easier to read.
Richard E. Barrans Jr., Ph.D.
Assistant Director
PG Research Foundation, Darien, Illinois
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