We have discussed several forms of energy - kinetic, potential, mechanical, thermal. Now
we will add another to our list: chemical energy. You utilize chemical energy whenever you use a
battery-operated device, burn fuel, eat food, or breathe oxygen.
The amount of energy stored in a molecule depends on the kinds of atoms in that
molecule as well as their arrangement with respect to each other. For example, we can pass an
electric current through water and break the water molecules into hydrogen and oxygen gas.
2H2O --> 2H2 + O2
We can experimentally determine that this process requires a net input of energy equal to
118 kcal. This is because a system containing two molecules of hydrogen gas and one molecule
of oxygen is more energetic (or less stable) than a system containing just two molecules of water.
Not only does the system with hydrogen and oxygen gas contain three molecules compared to
only two water molecules, a system containing molecules in the gas phase requires more energy
than a system containing molecules in a “cooler” liquid phase. The reverse reaction - the
formation of water from hydrogen and oxygen gas - will release 118 kcal of energy because that
reaction proceeds from a more energetic to a less energetic state.
Why is 118 kcal the magic number for water? What is the number for other substances?
By calculating exactly how much energy is required to decompose many different substances (as
well as other experiments), researchers began to build a model of how molecules are able to store
this energy. The result was the chemical bond model which states that the specific interactions
and associations among atoms within a molecule have specific bond energies associated with
them. In the case of water, the energy of each H-O bond is110 kcal, each H-H bond is 103 kcal,
and each O=O bond is 116 kcal. These numbers refer to the amount of energy required to break
the bond and also the amount of energy released when the bond is formed. Two hydrogen atoms
are much more energetic and less stable than one hydrogen gas molecule, so 103 kcal of energy is
released when the two hydrogen atoms combine to form one molecule. Thus the balance sheet for
the decomposition of water is
2H2O --> 2H2 + O2
2(110+110) --> 2(103) + 116
The total on the left side (the total input of energy required) is 440 kcal; the total on the
right side (the total release of energy) is 322 kcal. The difference of these two numbers is 118
kcal, the net input of energy required for the reaction. If there is a net release of energy, such as in
the reverse reaction, we call this free energy because it is available to do work. If you combined
hydrogen gas with oxygen gas, the result will be an explosion releasing 118 kcal of energy,
which, theoretically, one could harness to do work. In real life, however, the conversion of
chemical energy to work is very inefficient, as we will discuss later.
What exactly is a chemical bond? Nobody knows the answer to this question. For now,
the concept of chemical bonds is a model to help us quantitatively explain the energy difference
between a molecule and a set of free-floating atoms. We can develop the model in much more
detail, however.
First of all, it is important to know that there are two main kinds of chemical bonds: ionic
and covalent. Ionic bonds result from two atoms having very different electronegativities. You
can think of electronegativity as the ability of an atom in a molecule to attract electrons to itself.
See the table below, which is a modified version of the periodic table of elements:
http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/E/Electronegativity.html
The most electronegative element is fluorine (F) followed by oxygen (O). This
means that fluorine and oxygen atoms will attract electrons to themselves and away from
other atoms in the molecule. If two atoms having very different electronegativities, for
example fluorine and sodium (Na), interact, they will form an ionic bond. If the
electronegativities are similar, for example hydrogen (H) and carbon (C), they will form a
covalent bond, meaning that the electrons are shared equally between the two atoms.
If there is only a moderate difference in electronegativity, as occurs between
hydrogen and oxygen, a polar covalent bond may form. This is well illustrated in the
water molecule, where there are two H-O covalent bonds, but the electrons spend more
time around the oxygen atom than the hydrogen atoms, resulting in a polar molecule
having one positive end (or pole) and one negative end. If you have many water
molecules, then the positive pole of one water molecule will interact with the negative
pole of another molecule - this interaction is what is known as a hydrogen bond. I will
draw this on the board. The hydrogen bonding of water molecules is responsible for many
of water’s unusual behaviors that make it absolutely essential to life on Earth.
One more concept you should be aware of is that of double or triple bonds. A
single covalent bond involves the sharing of one pair of electrons. In the methane (natural
gas) molecule, for example, 4 hydrogen atoms share one pair of electrons each with one
carbon atom. This is the most stable arrangement of atoms because the carbon atom by
itself has only 4 electrons and needs 4 more electrons from other atoms in order to get to
its desired stable state of 8 electrons. Most of the atoms that we will be dealing with like
to have 8 total electrons (except hydrogen which only needs two electrons). In the carbon
dioxide molecule, carbon needs to get four extra electrons from just two oxygen atoms.
The only way this can occur is if each oxygen atom shares two pairs of electrons with
carbon so that each atom in the molecule ends up with 8 total electrons. Since two pairs of
electrons are being shared, each carbon-oxygen bond in carbon dioxide is called a double
bond and denoted with a double bar: C=O.
Below is a table of bond energies for common covalent bonds.
C-H
98 kcal/mole
O-H 110 kcal/mole
C-C
80 kcal/mole
C-O
78 kcal/mole
H-H 103 kcal/mole
C-N
65 kcal/mole
O=O 116 kcal/mole
C=O 187 kcal/mole
C=C 145 kcal/mole
Now we are ready to apply these concepts to larger issues we have been dealing
with in this class. For example, how much energy is released by burning natural gas,
gasoline, vegetable oil, or a candy bar? When we say “burning” we mean a reaction with
oxygen to convert the starting material into a lower energy state. A more technical term
for this is oxidation. If the reaction occurs quickly at high temperature, we call it
combustion. To calcuate energies of oxidation, we only need to obtain the chemical
formulas for the substances of interest, write the reaction describing their oxidation, and
then calcuate the bond energies on each side of the reaction. Keep in mind that the actual
energy obtained in the real world will differ significantly from our theoretical
calculations, but the main points remain the same.
natural gas (methane):
CH4 + 2O2 -> CO2 + 2H2O
4(98) + 2(116) -> 2(187) + 4(110)
624
814
net release of 190 kcal/mol
Gasoline consists of several different hydrocarbons (molecules that contain only
carbon and hydrogen) such as heptane, octane, nonane, and decane. Each of these
molecules only differs in the number of carbon atoms in the molecule. I will draw them
on the board. The reaction below is for octane.
2C8H18 + 25O2 -> 16CO2 + 18H2O
36(98) + 14(80) + 25(116) -> 32(187) + 36(110)
3528 + 1120 + 2900 -> 5984 + 3960
7548
->
9944
net release of 2396 kcal, 1198 kcal/mol
If we performed the same calculation for heptane, nonane, and decane, we would
see that bigger hydrocarbon molecules release more energy on combustion than smaller
molecules. Note that this value is per mole of octane, so it can be misleading to compare
directly this value to that of methane. If we express the energy released in terms of kcal
per gram, then octane releases 10.5kcal/g and methane releases 11.9kcal/g. One way to
think of this difference is that methane contains more carbon-hydrogen bonds per carbon
atom than octane. Thus it is the number of carbon-hydrogen bonds in a substance that is
the true currency of energy rather than simply the number of atoms.
We should pause for a moment to think about what this quantity of energy means.
If we were able to convert with perfect efficiency (not possible, but let’s continue
anyway) all 11.9 kcal/g of methane into electricity, then 1 kilogram of methane would
power a 60 watt light bulb for 10 days. That is a lot of energy packed into a small mass of
starting substance, and all of it is stored in the invisible bonds between atoms.
Let’s look at the combustion of some more substances. Ethane, C2H6, yields only
414 kcal per mole, since it contains only six carbon-hydrogen bonds per molecule. What
about ethanol, which many people mix with gasoline in car engines?
C2H5OH + 3O2 -> 2CO2 + 3H2O
5(98) + 80 + 78 +110 +3(116) -> 4(187) + 6(110)
1106
1408
net release of 302 kcal/mol
Ethanol releases less energy than ethane because, in effect, ethanol has already
been partially oxidized - it already contains an oxygen atom. Because of the energy
required to break the oxygen bonds in ethanol, it is less energy-rich than gasoline.
Precisely because it contains oxygen and few carbon atoms, though, ethanol burns cleaner
than gasoline, meaning it produces fewer nasty byproducts like carbon monoxide and
benzene.
Vegetable oils, when mixed with methanol, can be used as biodiesel. They are
composed of varying proportions of fatty acids, which burn, as with the hydrocarbons, to
yield carbon dioxide and water. The oxidation of palmitic acid, a major component of
many vegetable oils, is shown below.
C18H32O2 + 23O2 -> 16CO2 + 16H2O
31(98) + 15(80) +187 + 78 + 110 + 23(116) -> 32(187) + 32(110)
7281
9504
net release of 2223 kcal/mol or 7.94 kcal/g
When we eat food, we oxidize it with the oxygen we breathe to extract energy. We
write the chemical reaction for this process exactly the same way we have been writing
them for other fuels. For the sugar glucose, the reaction is:
C6H12O6 + 6O2 -> 6CO2 + 6H2O
7(78) + 5(110) + 7(98) + 5(80) +6(116) -> 12(187) + 12(110)
2878
3564
net release of 686 kcal/mol or 3.8 kcal/g
Glucose is the model of biological metabolic compounds - whenever you eat
carbohydrates, enzymes in your mouth and your stomach break complex carbohydrates
into simple sugar molecules, glucose being one of the most common. The simple sugar
molecules can then be transported through your bloodstream to individual cells where it is
oxidized - or burned - to provide energy for cellular processes.
Glucose is also the initial product of photosynthesis. In fact, the equation for
photosynthesis is exactly the reverse of the equation above for the oxidation of glucose.
The synthesis of glucose from carbon dioxide with a byproduct of oxygen is exactly what
green plants do, and this process requires an input of 686 kcal to produce one mole of
glucose. Where does this energy come from? The sun! Chloroplasts inside green plants
absorb sunlight and convert these photons into the chemical energy required to split water
molecules and combine the hydrogen atoms with carbon dioxide. We may discuss more
details of the biochemistry of photosynthesis another time. If you would like further
reading, I recommend
http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/
Just search for photosynthesis and browse the results. Much of the material from this
lecture came from this website.