Science Notes: Energy and Chemical Stability

Natural systems left to themselves move towards states of lower potential energy. For example, water flows down a hill or a ball rolls down a hill, if free to do so. States of lower potential energy are more stable. As a rule, the lower the potential energy of a system, the more stable it is. As a result, left to themselves, systems attempt to reach the configuration with the lowest energy possible under a given set of constraints. To change the state of a system from lower to higher potential energy, one must therefore supply energy to the system.

less stable
more stable
higher potential energy
lower potential energy
less bound
more bound

As more stable states have lower potential energy, we can get energy for use by moving a system to a lower potential energy. This is the basis for all energy transformation technologies.

Chemical Stability
Chemicals and their reactions are the medium through which nature stores and transforms energy. This energy is partly derived from the sun's pure electromagnetic energy that reaches the Earth as solar radiation, and partly from the energy stored in chemicals as potential energy in the chemical bonds. Recall the food chain diagram from the Introduction. Photosynthesis is an example of how nature stores and transforms energy via chemical bonds. Chemical bonds are essentially the phenomenon that atoms of elements stay close to each other, forming a compound, because that puts them in a lower (more stable) state of total energy.

These lower energy configurations of elements happen when the elements get an electronic configuration similar to the nearest inert gas. Electronic configurations of inert (noble) gases are the most stable in a given period (horizontal segment) of the Periodic Table. In fact, that is why these specific elements do not "need" to react with anything and are, therefore, chemically inert (or "noble"). These elements have no "need" to combine because their electron shells are completely filled with electrons. The noble gases are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Chemical activities of the elements are simplest to predict when they are close to the inert gases in the periodic table. Look at the first few periods of the Periodic Table in Figure 13. The activity can be described as a result of elements "wanting" to complete their shell. Although this is an anthropomorphic description, it is a helpful analogy.

Figure 13: First three periods in the Periodic Chart.
An interactive Periodic Table of the Elements, complete with date of element's
discovery, melting/boiling points, and electron configuration, is available at http://www.chemicalelements.com.

Lithium and sodium would tend to "lose" one electron to become more stable (more like the closest inert element). They can do this, for example, by combining with elements that can gain stability by adding one electron to their shell -- elements such as F and Cl. These pairs therefore favor ionic bonds, in which the electron is actually transferred.

When similar atoms such as H or O or N come together they can gain stability by sharing electrons in a covalent bond. Shared electrons spend time with either of the H atoms in H2 for equal amounts of time, for instance. In a case of a covalent bond such as those in H2O, however, the electrons spend more time on the average in the neighborhood of the oxygen.

The water molecule, which has an angle of 105° between the two H-O bonds, is therefore a polar molecule, being more negative at the oxygen end of the molecule, because the negatively charged electrons spend more time near the oxygen atom.

Nitrogen, hydrogen and oxygen atoms go to more stable configurations by forming the diatomic gases, N2, H2, and O2 respectively rather than remain in the atomic form: N, H, and O. When hydrogen atoms are produced in any reaction, pairs of these hydrogen atoms form covalent bonds with each other so that each has the helium (nearest inert gas) configuration at least a fraction of the time. Hydrogen has one electron and needs a total of two to be like He. So two hydrogen atoms share a pair of electrons, each belonging to one of the original atoms, thus forming H2. Schematically, it could be written as HxxH (x representing an electron). This schema is represented by H-H where the single line represents a bond consisting of two shared electrons.

Again, the same can be said for oxygen and nitrogen. Oxygen has an atomic number of 8, and has four electrons in the outermost shell. It needs two more to be like neon (nearest inert gas).

Exercise:

Nitrogen has __ electrons in the outer shell, needs __ to be more like ___

Carbon has __ electrons in the outer shell, needs __ to be more like ____

Oxygen has __ electrons in the outer shell, needs __ to be more like ____

Phosphorus has __ electrons in the outer shell, needs __ to be more like __

Hydrogen has __ electron in the outer shell, needs __ to be more like _____

answers

Following this logic, we can figure out the most frequent bond configurations for carbon, nitrogen, oxygen, phosphorus, hydrogen, with 4, 3, 2, 3 and 1 bonds respectively, as schematically in Figure 14.

Figure 14: Bond configurations for
carbon, nitrogen, oxygen, phosphorus, and hydrogen.

Energy transformations using chemical sources consist of changing the mutual configurations of these compounds accompanied by the release of energy, which we can then use for something. Chemical bonds "contain" energy that may be released when the bonds are made.


Representation of Bonds

Each atom has a number of bonds coming out of it equal to the number of electrons it shares in covalent bonds. So the line of a bond represents two electrons in activity, one from each of the two atoms it bonds. Thus for H2, H-H is really H:H, with each hydrogen atom contributing one electron to the bond.

Let's look at the example of nitrogen (atomic number = 7). Nitrogen, being 3 electrons short of its nearest inert gas (neon, atomic no. = 10), tries to bond so as to share 3 of it's electrons with other atoms bonding with it, in order to get a complete shell of 10 whenever possible. Thus it may share electrons with another N atom (forming N2), or with hydrogen (forming NH3).

Figure 15: Representations of compound using the bond scheme.

Similarly, carbon can bond with four hydrogens to form CH4 (a gas called methane), or with two oxygen atoms to form CO2 as depicted above. Note that we always have four bonds coming out of carbon, one out of hydrogen, and one out of oxygen. Look at their position in the periodic table to see why this is so. Atoms in the middle of the Periodic Table and their bonding become more complicated, and we will not need to deal with them here.

Carbon is the basis for all our life forms. It is a very versatile atom because of its capability to form four bonds. Depending on the amount of hydrogen available to bond, and the temperature and pressure conditions, carbon can form a plethora of compounds with hydrogen alone. One such family is the hydrocarbons, important in our context because they are the basis of fossil fuels.

Note how some of these compounds have double and triple bonds between carbons. This happens when carbon and hydrogen combine under circumstances in which there is not enough hydrogen to satisfy all four bonds of each carbon. For example, if there is plenty of hydrogen to combine with carbon, we get CH4 or C2H6 (Ethane), with all single bonds. With less hydrogen we get C2H4 or C2H2 (less hydrogen for the same number of C atoms). C2H4 has a double bond between the carbons, and C2H2 has a triple bond. Compounds with double and triple bonds are called unsaturated, while single bond compounds like C2H6 are said to be saturated. Unsaturated compounds are more reactive than saturated compounds because not all the C atoms are bonded to four other atoms. Hydrocarbons are not the only compounds that can be unsaturated. Carbon monoxide is a good example of an unsaturated compound, "looking" for another oxygen atom to form CO2, a more saturated compound. When carbon (in coal or wood, for example) burns in an environment with insufficient oxygen, it forms CO which is deadly when breathed in. This is the reason to ensure plenty of access to fresh air when we have a fireplace or running car engine.

Note that many representations are two-dimensional, and that in actuality, the electrons forming the bonds are distributed in three dimensions. In a compound like CH4, the carbon is in the middle of a tetrahedron with the 4 H atoms at the vectors.

 

 

Figure 16: Linear hydrocarbons.

Similarly there can also be C3H8 (propane), C4H10 (butane), C5H12 (pentane) and so on. When a formula is written as CH4 just showing the proportion of atoms, it is called an empirical formula. When the bonds are shown as in Figure 15, it is called a structural formula. A single empirical formula may represent different compounds because the structures may be different for the same number of atoms combining.

Try drawing propane, butane, and pentane. Note that there are always four bonds coming from carbon. The linear structures are called aliphatic hydrocarbons. In addition to the linear hydrocarbon molecules, hydrocarbons may also be formed into ring structures. The ring structure possesses the property that enables us to smell these compounds! So they are called aromatic hydrocarbons. The simplest aromatic hydrocarbon is C6H6, benzene. The structure of benzene was long a puzzle in chemistry, with chemists wondering what the structural formula for C6H6 could be. They knew the empirical formula was C6H6. It is said that the great organic chemist Kekulé, who had been wondering about this, dreamed one night of a snake swallowing its tail and was inspired to draw the ring structure! Note the alternating single and double bonds, a clever way of ensuring four bonds from each carbon shell.

The versatility of carbon in forming bonds, ring structures and various configurations is the basis of life on our planet. The chemistry of carbon compounds is therefore called organic chemistry. More complicated carbons compounds are described in the Ecological System and Materials System. For now, let us look at some additional aromatic and aliphatic compounds, and note some aspects that are relevant to energy storage and release.

Aliphatic hydrocarbons are the basis of fossil fuels. All saturated hydrocarbons react with oxygen at high temperatures to form carbon dioxide and water, and give off energy. This oxidation reaction is the basis of the internal combustion engine. Gasoline normally contains hydrocarbons from C6 to C18, a mixture of over 100 compounds! An example reaction of the combustion of a hydrocarbon is:

C7H16 + 11O2 7CO2 + 8H2O + energy

"Burning" (or a combustion reaction) consists of combining with oxygen at high temperatures. The combustion reaction of acetylene (C2H2) with oxygen gives off such a large amount of energy that it is used as a welder's torch.

Ring compounds do not play as large a role in energy production but often occur as byproducts or waste products. These polyaromatic hydrocarbons (PAH's) pose a serious pollution problem.

Ring compounds, based on the benzene ring, are so common in biochemistry that we just draw to represent C6H6. Adding one more carbon and two hydrogens to the benzene ring gives us C7H8 which is methyl benzene or toluene (at right).

Ring compounds can get very complicated. Several organic compounds playing an important role in our physiology are shown in the Ecological System.

Chemical Energy Release and Bond Energies
The amount of energy released when a bond is formed between atoms is called the bond energy. Bond energies represent a state of potential chemical energy. We can get energy from a system as it moves from a state of higher potential energy to one of lower potential energy (e.g. water falling). Chemical reactions in which the compounds formed after a reaction (called products) have lower total bond energy than the reactants can release chemical energy. Such reactions in which energy is given off are called exothermic (or more correctly, exoergic) reactions. Conversely, reactions that absorb energy are said to be endothermic.

Table 8 gives the energies for bonds we will commonly encounter. The table defines these energies in units of kcal/mole. A mole is an abbreviation for "gram-molecular weight" of a substance.

BOND
Energy
(kcal/mole)
H-H
104
C-H
99
C-C
83
C=C
146
CC
200
O-O
35
O=O
119
O-H
111
C-O
86
C=O
177
 
BOND
Energy
(kcal/mole)
H-F
135
H-Cl
103
H-Br
87
H-I
71
N-N
39
NN
225
N-H
93
Cl-Cl
58
Br-Br
46
I-I
35
Table 8: Bond Energies.
(the bond energy is expressed in kcal/mole.)

Let us see what the energy values in Table 8 mean. The bond energy of H-H is 104 kcal/mole. This means that when hydrogen atoms combine to form molecular hydrogen H2, represented by the reaction H + H H2, for every mole (2g) of H2 formed, 104 kcal of energy are released. Conversely it takes 104 kcal to break apart a mole (6 x 1023) of hydrogen molecules. From this we can draw a simple chemical energy level diagram for the above reactions, analogous to Figure 12 of gravitational potential energy.

Figure 17: Energy Level Diagram of H2.

One mole of an H2 (2g) molecule has 104 kcal total potential energy less than 2 g of H atoms. This is why when H atoms are formed in a reaction, and these atoms are the only atoms available, they combine to form H2 (roll down the potential energy "hill" towards a more stable state). In forming the H-H bond, 104 kcal of energy are released for every 2 g (6.02 x 1023 molecules) of hydrogen gas (H2) formed. Similarly oxygen, and nitrogen exist as O2 and N2 rather than in the atomic form as O and N. So whenever we say hydrogen, oxygen, or nitrogen gas, we mean H2, O2, N2. For H, O, N we specifically say atomic hydrogen, atomic oxygen, and atomic nitrogen.

 

 

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