Why is a knowledge of basic chemistry essential to understanding biological principles? Chemistry is a study of the basic, non-living portions of our physical world. Biology is the study of the living. The relationship between these two apparently opposite disciplines has only been emphasized heavily in the last thirty years. Man, in the search for the roots of his existence, has come to believe that ultimately the mechanics of life are explicable in chemical and physical terms. The "whys" of life may never be explained, but the mechanics may be entirely explained in terms of specific and definable chemical interactions. Biology is no longer a science consisting mainly of cataloging and categorizing life forms. Today's emphasis and thrust is in a relatively new direction -- Biochemistry -- the interactions of the unique chemical reactions that are expressed in the miracle we call life. What we are and what we eat can be reduced to chemicals.
 
SUBJECTS COVERED BY THIS REVIEW:

CHEMICAL ELEMENTS
ATOMS
CHEMICAL COMPOUNDS
THE IONIC BOND
THE COVALENT BOND
OXIDATION STATES
CHEMICAL REACTIONS
OXIDATION-REDUCTION REACTIONS

At present there are ninety naturally occurring chemical elements and several that have been artificially created. Each element has its own distinct and definable properties, and the relationships between the elements is summarized in a chart called the periodic table, which groups the elements according to similarities in their chemical properties. For convenience each element is represented by a symbol.

Some of the most biologically important elements and their symbols are as follows:
 

Although there are similarities between elements, each element is a distinct entity. The smallest particle that retains the chemical properties of an element is an atom. Atoms of all elements are composed of the same subatomic particles, but it is the numbers of these particles and the balances between them that make an atom of an element distinct from the atoms of each of the other elements.

PERIODIC TABLE

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As stated previously, the smallest particle that retains the chemical properties of an element is an atom.  Each atom is composed of three types of subatomic particles:  electrons, carrying a negative electrical charge; protons, carrying a positive electrical charge; and neutrons, carrying a neutral or zero electrical charge. It is the numbers of these particles and the balances between them that gives an atom the distinct chemical properties of each of the different elements.

Several models of atomic structure have been created and then updated as more has been learned about atomic structure. Although not the most modern model of atomic structure, the best model for easily and simplistically presenting basic chemical principles is called the Bohr Model, named after its discoverer. This model was quite popular for many years and will suffice for a generalization of atomic structure.

Bohr visualized the structure of the atom in a form very similar to the structure of our solar system. Our solar system consists of a center, the sun, with planets revolving around that center in definite paths or orbits. Bohr visualized the structure of an atom as having a center, called the nucleus, composed of protons and neutrons with electrons revolving around that nucleus in defined orbits or shells. Unlike our solar system which contains only one planet in each orbit or path around the sun, electron orbits or shells have capabilities of containing more than one electron depending largely on how close to the nucleus the orbit is and how broad a circle the orbit forms. For example, the orbit or path of electrons closest to the nucleus is capable of holding two electrons. The orbit next closest to the nucleus can hold as many as eight electrons; and the third closest orbit can hold eight electrons. Here we will not consider any elements having enough electrons to exceed filling the first three orbits. It should be noted that depending on the number of electrons in an atom, the orbits will fill from inside out until the maximum number of electrons for a given atom is exhausted. Let's visualize an atom having three orbits, each orbit filled with the maximum number of electrons.
 
 

Here an atom with enough electrons (18) to fill three orbits or shells is represented. It would be the element argon, but how would one know that it is argon being represented here?

The structure of an atom of any given element can be determined by knowing two pieces of information:

1. atomic number - the number of protons in an atom of each of the elements. This number can be determined by looking at a periodic table which lists the atomic numbers of the elements. The atomic numbers of the elements for which the symbols were listed previously are as follows:

2. balance of charges - The Bohr model of atomic structure dictates that a normal atom be electrically neutral overall. Remembering that atoms are composed of three types of subatomic particles: positive (+) protons, negative (-) electrons, and neutral (0) neutrons, it can be deduced that for the atom to be neutral overall the positive and negative charges of the subatomic particles must balance. The number of protons is equaled by the number of electrons.

• The smallest part of an element is an atom.
• An atom is composed of subatomic particles:  (+) protons, (-) electrons, and (0) neutrons.
• Protons and neutrons form the nucleus of an atom.
• Electrons revolve around the nucleus in defined circular paths or orbits.
• Orbits have a maximum number of electrons they can contain:  two in the first, eight in the second, and eight in the third.
• Electrons fill the orbits from inside out.  The number of orbits and the fullness of the outermost orbit depend on the number of electrons an atom has.
• Atoms are electrically neutral; (+) protons balance (-) electrons.
• The atomic number of an element indicates the number of protons in an atom.

Using all the preceding information, let's look at the structure models of some of the elements.
 

Ex.1 Hydrogen (H)   atomic number one (1)  Hydrogen has one proton in its nucleus. The proton is balanced by one electron. The one electron goes in the first available orbit. [Note that this would not fill up this orbit.]

 
 

Ex.2 Carbon (C)   atomic number six (6)  Carbon has six protons. The protons are balanced by six electrons. Orbits fill from inside out with two electrons going into the first orbit and four electrons going into the second orbit. [Note that the second orbit would not be full, as it is capable of containing eight electrons.]

 
 

Ex.3 Oxygen (O)  atomic number   eight (8)  outer orbit not filled	Ex.4 Sodium (Na)   atomic number   eleven (11)  outer orbit not filled	Ex.5 Chlorine   atomic number   seventeen (17)   outer orbit not filled

 
Having determined the general structural pattern for atoms, we can now begin to approach the concepts of why atoms of different elements react with each other or don't react with each other; and, in addition, how atoms react with each other.

The why in the reactivity of atoms can be illustrated by a natural law of Physics. In simple terms this law states that any system left to itself will move toward its greatest state of stability. As an example of this law, visualize a marble placed on a board that is situated at an angle.

When the marble is released, gravity will cause the marble to roll down the board to an area that is level; and then the marble will stop. Having the marble on an inclined plane is an unstable system. Release the marble and leave it alone, and it will move to its greatest state of stability.
 
Now, let's apply this natural law to a chemical system. The two simplest elements hydrogen and helium have very similar physical properties.  As we will discover, their chemical properties are vastly different. Hydrogen is atomic number one and helium is atomic number two. Both hydrogen and helium are lighter than air gases. Hydrogen is lighter than helium and was used extensively in lighter than air aircraft in the early part of this century. However, this practice was abandoned, for, although hydrogen was very effective, it is tremendously reactive resulting in explosions that are numbered among our worst air disasters. Today, helium, although not as light as hydrogen, is used. Helium is non-reactive--or in chemical terms--inert, and is safe enough to be used in balloons for children. What is the difference between these two gases that makes one of them highly reactive and the other inert? Let's look at models of their atomic structure.
 

	Hydrogen (H)  atomic number one (1)  There is one electron in the outermost orbit of a hydrogen atom. That orbit could contain two electrons.    Helium (He) atomic number two (2)  There are two electrons in the outermost orbit of a helium atom. The orbit contains its maximum number of electrons.
It is because helium atoms are at maximum stability with full electron shells and balanced electrical charges that they are stable and unreactive. Hydrogen atoms have balanced electrical charges but do not have full outer shells and are therefore reactive.

 
 

Left to itself, any system will move toward its greatest state of stability. As with hydrogen atoms, atoms lacking full outer shells will move toward filling their outer shells with electrons. They do so by reacting with other atoms forming chemical bonds, resulting in the creation of chemical compounds.

 

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A chemical compound is two or more atoms bonded together. A chemical compound is represented by a chemical formula. The chemical formula tells what kinds of elements are involved in a compound and how many atoms of each element are involved in forming the compound.

Note the following examples of compound formulas.
 

H2O	This represents one unit of the compound water. The formula indicates that in one unit of water there is one atom of oxygen (O) and there are two atoms of hydrogen (H).
2 H2O	This simply represents two units of the compound water, and it does not affect the relationship between the hydrogens and the oxygen that are within the unit. In other words, this is not the same as H4O2.
CH4	This represents one unit of the compound methane. Each unit is composed of one carbon atom and four hydrogen atoms.
CH3OH	This is methyl alcohol--one carbon, one oxygen, and four hydrogens. There is a reason that it is not written as CH4O, but this need not concern us here. CH3OH and CH4O are not the same compound.
H2	This is the compound hydrogen rather than an atom of hydrogen. The compound hydrogen is composed of two atoms of hydrogen chemically bonded together.

 
Compounds are formed by atoms chemically bonding together. There are two major types of chemical bonds -- the ionic bond and the covalent bond. In both cases these bonds between atoms are the result of atoms trying to achieve greater stability by attempting to fill their outer electron orbits.

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Some atoms attempt to achieve full outer electron orbits by transferring electrons into their outermost shells, or out of their outermost shells. As an example, let's look at sodium atoms and chlorine atoms.
 

Sodium   atomic number eleven (11)   It has one electron in an outermost shell that could hold eight. If it could get rid of that electron, the shell below would become its outermost shell; and it is filled.	Chlorine   atomic number seventeen (17)  It has seven electrons in an outermost shell and only needs one more electron to fill it.
When sodium atoms and chlorine atoms are close together, an electron will transfer from sodium to chlorine. This gives both sodium and chlorine full outer electron shells.

 
Note that after this transfer occurs the sodium and chlorine atoms no longer have an internal electrical balance. Sodium, having lost an electron, has become positive by a factor of one and becomes the sodium ion (Na⁺). The chlorine, by receiving an extra electron, becomes negative (electrons are negative particles) by a factor of one and is known as the chloride ion (Cl^-). By transferring an electron, sodium and chlorine have become more stable by acquiring full electron shells; but they have become more chemically reactive by the unbalancing of their electrical charges. Once again, any system left to itself will move toward its greatest state of stability. The sodium ion (Na⁺) and the chloride ion (Cl^-) will be attracted to each other. These ions will stick together to balance out their electrical charges and form the compound sodium chloride (Na⁺Cl^-). The overall charge on a unit of sodium chloride is zero, and the compound is stable. [Sodium chloride is table salt.] The ionic bond can be defined as the attractive force between positively and negatively charged ions. In addition, elements, like sodium, that tend to lose electrons are called metals; and those, like chlorine, that tend to gain electrons are called non-metals.

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Most atoms tend to share electrons between their outer shells instead of actually transferring them. When this occurs, a covalent or molecular bond is formed. A covalent bond can be defined as one pair (2) of shared electrons between atoms. Let's examine some compounds having covalent bonding.
 

The compound methane (CH4) is a covalent compound. In methane one carbon atom forms four different covalent bonds (shared pairs of electrons) with four different hydrogen atoms.
	Carbon   atomic number  six (6)  Carbon needs four electrons to fill its outer shell.	Hydrogen   atomic number one (1) Hydrogen needs one electron to fill its outer shell.
If four hydrogens share their electron with carbon and carbon shares each of its four electrons with hydrogen, all will fill their outer shells. This sharing also results in the formation of the covalent compound methane.  We can represent the covalent bonds in methane by showing all the electrons and the sharing of electrons. Four covalent bonds are formed. Each bond involves one electron from carbon and one electron from hydrogen--a pair of shared electrons.

To represent the bonding in methane, it is not actually necessary to show the nucleus of the atoms or all the shells of the atoms. For simplification, electron dot notation can be used to show the bonding between the carbon and the hydrogens.
 

Only the shared electron pairs are shown between the symbols for the elements involved. Remember that in each case the pairs of dots represent one carbon electron and one hydrogen electron -- a shared pair of electrons; a covalent bond.
To further simplify this representation, line notation is often used. A line is used to represent the shared pairs of electrons rather than two dots.

Let's examine one other example of a covalently bonded molecule--a molecule of oxygen (O₂).
 

Oxygen  atomic number eight (8)   An atom of oxygen needs two electrons to complete its outer shell. This can be accomplished if two atoms of oxygen share electrons with each other.

 
 

In electron dot notation this would be shown as:    In line notation this would be shown as:

Having discussed the major types of bonding and chemical formulas, we can now center our attention on a single common, but somewhat chemically amazing, compound--water. Water is considered to be the universal solvent because so many different compounds will dissolve easily in water.  A solvent is a substance in which other substances will dissolve. Dissolved substances are called solutes. The combination of solvent and solute is a solution.
 

The formula for water is most often written as H2O but is more properly written as HOH. Water exists in both a covalent and an ionic state. In its ionic state it is composed of the hydrogen ion (H+) and the hydroxyl ion (OH-). [There is a covalent bond between the oxygen and the hydrogen in the hydroxide ion.] In its covalent form water is more properly represented as pictured at left.  Each hydrogen is sharing with the oxygen, but the sharing is not equal. Oxygen is such a strong electron puller that in a water molecule it is slightly negative and the hydrogens are slightly positive. This makes water a polar molecule. It has a negative pole and a positive pole. Because of the dual nature of both ionic and covalent water molecules, water exerts a strong electrical pull on substances placed in it dissolving many of them easily.

 

Water also demonstrates a third type of bonding called hydrogen bonding. A hydrogen bond is a weak attractive force between a slightly positive hydrogen and a strongly, electrically negative atom such as oxygen or nitrogen. Consider three covalent water molecules as an example.  Because water molecules are polar, there is an attraction between the slightly positive hydrogen pole of one water molecule and the slightly negative pole of the other water molecule. This is a hydrogen bond.  While the hydrogen bond is very weak, hydrogen bonding can exert a tremendous force due to the large number of water molecules in any aqueous solution.

 
 
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Consider, for example, the sodium and chlorine atoms becoming ions. Sodium atom with a zero electrical charge loses an electron and becomes the sodium ion with a plus one electrical charge. The oxidation state of the sodium ion is +1 indicating that one electron is displaced away from the positive nucleus. The chlorine ion gains an electron giving it an oxidation state of -1 indicating that one electron is displaced toward its nucleus overbalancing the charge of the nucleus by a factor of one. By losing an electron, sodium is said to have been oxidized. By gaining an electron, chlorine is said to have been reduced.

In covalent compounds electrons are not lost or gained but are shared instead. However, the atoms in covalent compounds still have oxidation states because some atoms are stronger electron pullers as they share the electrons of other atoms. Pulling extra electrons would result in a negative oxidation state, and atoms that have their electrons pulled achieve a positive oxidation state. Which atoms pull electrons in a compound or have their electrons pulled is situational and varies from compound to compound. However, some atoms are fairly consistent. Oxygen tends most often to pull the two electrons it needs and have an oxidation state of -2. Hydrogen tends to have its electron pulled and have an oxidation state of +1. When two like atoms are sharing, neither tends to out-pull the other; and both have an oxidation state of zero (0). Let's look at the oxidation states in a molecule of water.

The oxygen pulls an electron from each hydrogen giving it an oxidation state of -2. Each hydrogen has an oxidation state of +1. Note that all the oxidation states in the molecule add up to zero and that the molecule has a negative and a positive end.

Consider the molecule of oxygen.  Four electrons are being shared, but neither oxygen is stronger than the other. Each oxygen has an oxidation state of zero. Note that the oxidation states add up to zero.

Now consider methane. Each hydrogen has an oxidation state of +1, following the rule that hydrogen tends to have its electron pulled by other atoms; but we have no rule for the carbon. However, we do know that the oxidation states should add up to zero. Therefore the oxidation state of the carbon must be -4.
 
 

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A chemical reaction is basically a change in the chemical state of atoms whether they be free atoms or atoms already involved in forming compounds. Reactions generally fall into one of three types.

1.    Synthesis reactions

A + B ® AB
Substance A combines with B to yield AB.

AB ® A + B
Substance AB decomposes to yield substances A and B.

AB + CD ® AD + CB
Substances AB and CD exchange to yield substances AD and CB.

Dissociation plays roles in living processes other than electrolyte balances. The sources of the terms acid and base lie in the process of dissociation. For our purposes here we can define an acid as any compound which dissociates releasing the hydrogen ion (H⁺) into solution. A base can be defined as any compound which dissociates releasing the hydroxide ion (OH^-) into solution.  Note that you may be familiar with more modern and more accurate definitions for acids and bases.

 
Hydrochloric acid dissociates to yield Hydrogen and Chloride ions.
 

Sulfuric Acid dissociates to yield Hydrogen and Sulfate ions.
 

Sodium Hydroxide dissociates to yield Sodium and Hydroxide ions.
 

Potassium Hydroxide dissociates to yield Potassium and Hydroxide ions.

These are just a few examples of the many different acids and bases. Acids and bases differ in strength, but their differences in strength are due simply to how much they dissociate to release hydrogen ions or hydroxide ions. It is the amount of hydrogen ion or hydroxyl ion that determines how caustic an acid or base may be. Strong acids release (dissociate) large amounts of hydrogen ion; weak acids release (dissociate) small amounts of hydrogen ion. The same is true of the hydroxide ion of bases.

The strength of an acid or base solution is measured on a scale called the pH scale. The "p" stands for power or concentration; the "H" stands for hydrogen ion. The scale has a range from 0 to 14, and the numbers represent the negative logarithm of the concentration of hydrogen ion in the solution being measured. A pH of 7 is neutral and indicates the concentrations of hydrogen ion and hydroxide ion are equal. Water, when pure, has a pH of 7 because dissociated water releases equal amounts of hydrogen and hydroxide ions.

The proper pH for most living cells ranges between 6 and 8 depending on the type of cell. In most cells this range is maintained by ions known as buffers. A buffer ion is an ion which controls the pH between certain limits by combining with or releasing hydrogen ions as needed.

One of the most common buffering ions in living systems is the bicarbonate ion (HCO₃^-1). The bicarbonate ion combines with hydrogen ion when it is in excess and will release hydrogen ion when it is in low supply.

Before leaving the topic of acids and bases, one additional property should be noted. When an acid and a base are combined, an exchange reaction will occur producing water and a salt.

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The compounds we consume for energy are high energy carbon compounds. The original source of the energy in these compounds is sunlight and the process of photosynthesis. Cells in our bodies by a series of exchange reactions convert the high energy compounds into low energy carbon compounds which we release as a waste product. The reactions that accomplish this energy conversion are called oxidation-reduction reactions.
 
 

Oxidation can be defined as the loss of electrons by an atom. Reduction can be defined as the gaining of electrons by an atom. Put even more simply, oxidation is electron loss; and reduction is electron gain. Let's examine a series of oxidation-reduction reactions as might be accomplished by a cell to acquire energy.

There is a simple memory tool you can use:  LEO the lion says GER.  LEO:  Loss of electrons is oxidation.  GER:  Gain of electrons is reduction.

  Examine the oxidation states of the atoms on the left side of the equation as compared to those on the right side. You should notice two types of atoms changed oxidation states because of the reaction. Carbon changes from -4 to a -2. This is oxidation. Two electrons have shifted away from the carbon toward an oxygen (electron loss). Hydrogen has changed from +1 to a 0 oxidation state (electron gain).
  Carbon changes from -2 to 0 (oxidation).
Oxygen changes from 0 to -2 (reduction).
  Carbon changes from 0 to +4 (oxidation).
Oxygen changes from 0 to -2 (reduction).

In the section in this chapter that describes the ionic bond we saw that sodium atoms tend to lose electrons to chlorine atoms creating the sodium and chlorine atoms. This is another example of oxidation and reduction. Sodium loses an electron and is thus oxidized. Chlorine gains an electron and is thus reduced. Can you explain how this is slightly different from the other examples in this section?

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