II. The Chemical Basis of Life - Subatomic, Atomic, and Molecular

Levels of Organization

All living systems are essentially complex systems of chemical reactions. Consequently it is essential to understand the chemical basis of life if one is to understand its organization and functioning at the higher, more visible levels. In this section we will explore the first three levels of complexity which constitute basic chemistry.

A. Elements and atoms - All of the matter of the universe is made up of 92 different types of atoms or elements.

1. Element - Any substance which cannot be further subdivided by chemical means. An element contains only one kind of atom. Examples include such things at oxygen, hydrogen, sulfur, lead, iron, gold, carbon, etc.

2. Atom - The smallest part of an element which still retains the chemical and physical properties of that element.

B. Atomic structure - Atoms are composed of three subatomic particles known as protons, neutrons, and electrons.

1. Protons and neutrons form the dense center or nucleus of the atom and the electrons orbit around this nucleus.

2. Protons are electrically positive and electrons are electrically negative. Neutrons have no electrical charge and are therefore neutral.

3. For an atom to be electrically neutral, it must have the same number of protons and electrons.

C. Atomic number - This is the number of protons in the nucleus. Each succeeding element has one more proton. The simplest element (atom) is hydrogen which contains one proton and has an atomic number of one. Uranium has an atomic number of 92 as it contains 92 protons.

D. Atomic weight - The weight of each atom is obtained by comparing it to the weight of carbon which is assigned a weight of exactly l2.0000. The atomic weight of any atom is approximately equal to the numbers of protons and neutrons found in the nucleus and is termed the mass number. Atomic weight is measure in Daltons (atomic mass units).

E. Molecule - This is two or more atoms which have combined chemically. If the atoms are of different types then the molecule is referred to as a compound.

F. Molecular weight - This is the weight of a given molecule. It is equal to the sum of the atomic weights of all of the atoms found in that molecule.

Example - Water - H20 - The atomic weight of hydrogen is l and that of oxygen is l6. Therefore water is l + l + l6 = l8.

Biological molecules are often very large. A single protein molecule may have a molecular weight of 500,000, while a molecule of DNA can reach into the billions.

G. Chemical bonds - The atoms which make up molecules are held together by energy interactions involving the electrons. Such electron interactions are known as bonds.

1. Covalent bonds - This is the type of bond formed when two atoms share one or more pairs of electrons.

a. Electrons exist in energy levels or shells, around the nucleus. Each shell can hold a finite number of electrons. The first level or shell is filled when it contains two electrons. The second shell and each succeeding shell is filled when eight electrons are present.

b. An atom reaches maximum stability when the outer shell is filled with electrons. Atoms with unfilled outer shells can effectively complete these shells by sharing electrons with other unfilled atoms, and this is the basis of covalent bonding.

c. An example of covalent bonding is water. Oxygen has six electrons it its outer shell and therefore requires two more to become stable. Hydrogen has one electron in its single level and requires a second electron for filling of this shell. Two hydrogen atoms sharing with one oxygen atom effectively fills all of the outer shells and two covalent bonds are formed in the process bonding of hydrogen to oxygen.

(1) Non-polar covalent bonds - When two atoms share a pair of electrons equally between them the result is the formation of a non-polar covalent bond. Such a bond is electrically neutral.

(2) Polar covalent bonds - When two atoms share a pair of electrons unequally between them. In this case, the pair of electrons is pulled closer to the nucleus of one atom than the other. The result is that one atom becomes more negative that the other and electrical polarity (positive and negative poles) develop.

(a) Water is an example of a polar covalently bonded compound. The oxygen atom pulls the electrons of the hydrogen closer to its nucleus than that of the hydrogen nuclei. The result is that each hydrogen has a positive proton sticking out and therefore moves toward positive polarity while the oxygen end with two extra electrons moves toward negative polarity.

2. Hydrogen bonds - These are weak electrical attractions formed between hydrogen (which tends to be positive in many covalent molecules for the reasons stated above) and certain very electronegative atoms, oxygen and nitrogen to be specific. These bonds are very weak and easily disrupted, but they play very important roles in biological chemistry.

    1. Water is a good example of the importance of hydrogen bonding. The weak hydrogen bonds that form between the hydrogen atoms in one molecule and the oxygen atom in another molecule result in water molecules being held together in close association. The practical results of this is the fact that water, without hydrogen bonding, would be a gas at body temperature, and would not become a liquid until well below 0o C.

H. Ionization and ionic bonds - Ionization is that process by which a charged atom or ion is formed. Ions form whenever an atom gains or loses one or more electrons.

1. Examples: Sodium loses an electron and becomes positive (Na+).

Chlorine gains an electron and becomes negative (Cl-).

Calcium loses two electrons to become positive (Ca++).

2. Ionic bonds - This is a type of bond which is formed when one atom pulls an electron away from another atom. The atom gaining the electron becomes negative while the atom losing the electron becomes positive. The positive and negative charges which have developed attract one another and hold the atoms together. There is no sharing of electrons. An electron is actually transferred from one atom to another. In their solid state ionic compounds form crystals known as salts.

a. Example: Sodium chloride (NaCl). Sodium has one electron in its outer shell. Chlorine has seven in its outer shell. Sodium loses one electron to chlorine. Now both have their outer shells filled and are stable. Each has also become charged, and the unlike charges hold the ions together.

I. Dissociation and electrolytes

1. Ionic compounds come apart or dissociate when dissolved in water. The water molecules move between the ions which then move freely about in the water.

2. Covalent compounds with polar bonds may also dissociate in water but are not as likely to do so as ionic compounds because the covalent bond is stronger than the ionic bond.

3. Compounds which dissociate to form ions are termed electrolytes because they will conduct electricity in solution.

4. Electrolytes are important to many physiological processes. These include (but are not limited to) nerve conduction, muscle contraction, and energy production.

5. It is important to distinguish dissociation whereby molecules actually come apart from dissolving in which water molecules move between the intact molecules. Sugar dissolves in water but does not dissociate. Salt dissociates in water.

J. pH - This is defined as the negative logarithm of the hydrogen ion concentration of a solution. It is usually written in mathematical form as

pH = 1/log (H+)

l. pH is measured on a scale of 0 - l4. pH 7 is said to be neutral while values less than 7 are acid while those greater than 7 are basic or alkaline.

2. pH is very important because all biochemical reactions are pH sensitive. Each reaction has its own optimum pH value at which it occurs best and if the pH value deviates in either direction from this optimum value the reaction begins to slow down and eventually ceases completely.

3. A good example of the critical nature of pH is the pH values of the blood listed below.

Normal - 7.4

Acidosis - 7.2 or less

Alkalosis - 7.6 or more

Coma and death occur at 6.8

K. Acids and bases - These are substances that will alter pH.

1. Acid - Any substance that will yield a hydrogen ion when dissolved in water.

2. Base - Any substance which will remove a hydrogen ion from solution when dissolved in water.

3. Strong acids or bases are those which will ionize almost completely when placed in water. They will either add or remove large amounts of hydrogen ion thereby effecting pH significantly.

4. Weak acids or bases are those which only partially ionize in water. They do not add or remove large amounts of hydrogen ion and therefore do not effect pH nearly as much as equivalent amounts of strong acids or bases.

5. Acids and bases react with one another to form salts. For example:

HCL + NaOH = NaCl + HOH

6. Acids and bases are physiologically important because they alter the pH of the body fluids. Salts are significant because they represent major sources of essential electrolytes.

L. Buffers - A buffer is any substance which when added to a solution resists a change in pH. Buffers work by converting strong acids and/or bases into weak acids and/or bases.

1. Sodium bicarbonate is the major buffer found in the blood and tissue fluids. It functions in the following way.

NaHC03 = Na+ + HCO3-

HCl (strong acid) = H+ + Cl-

NaCl (salt) + H2C03 (weak acid)

The weak acid (carbonic acid) formed can now react with strong bases like NaOH into the weak base sodium bicarbonate.

NaOH = Na+ + OH-

H2CO3 = H+ + HC03-

HOH (water) + NaHCO3 (weak base)

M. Isotopes - Atoms which have identical chemical properties but differing atomic weights.

1. The weights are different because of differences in the number of neutrons. The isotopic form which occurs most abundantly in nature is considered the "normal" form. Isotope forms are designated by the elements symbol and a superscript indicating the atomic weight. For example, in the case of carbon, we find Cll, Cl2, Cl3, and Cl4.

2. Many of the rare isotopes are unstable and undergo nuclear reactions to form stable forms. In doing so they often emit radiation (radioactivity).

3. Isotopes are important in several areas.

a. Research - Radioactive isotopes can be incorporated into molecules and the fate of the isotope then traced through a series of reactions. Most of our knowledge of biochemistry has been derived from these types of isotope studies.

b. Therapy - Certain isotopes (Cobalt 60 for example) produce very high energy radiation that can be used to treat cancer.

c. Diagnosis - Radioactive isotopes can be injected into the body and then traced. They indicate the chemical activity of different tissues.

N. Solutions

1. Definition - A homogeneous mixture of two or more components which cannot be distinguished and which do not settle out.

2. The solvent is the component in greatest amount while the solute is the component in the lesser amount. In biological systems water is always the solvent. Water is a polar solvent with a number of unique properties which make life as we know it possible.

3. Expressing concentrations of solutions - It is frequently necessary to know the concentration of a solution. Generally concentrations are expressed for the solutes although it is also correct to express solutions in terms of solvent concentrations. There are several different ways of expressing concentrations, but only two will be considered here.

a. Percentage - This is the percent of solute in the solution. It may be expressed in several ways.

(1) Wt./Wt. - 10% solution = 10 grams of solute plus enough solvent to make 100 grams of solution.

(2) Vol./Vol. - 10% solution = 10 ml of solute plus enough solvent to make 100 ml of solution.

(3) Wt./Vol. - 10% solution - 10 grams of solute plus enough solvent to make l00 ml of solution.

As one ml of water weighs one gram, Wt./Wt. solutions and Wt./Vol. solutions become one and the same when water is the solvent.

b. Molar solutions - This solution is based upon a unit known as a mole. A mole is equal to a gram-formula-weight which is the molecular weight of a molecule in grams.

(1) A l.0 M solution = One mole of solute plus enough solvent to make one liter of solution.

A 0.l M solution = 0.l mole of solute plus enough solvent to make one liter of solution.

(2) The advantage of using molarity as a measure of concentration is that all solutions of equal molarity have the same number of solute molecules.

(3) Example - Prepare a one molar solution of NaCl. The molecular weight of NaCL is 58. Therefore a mole or gram-formula-weight of NaCl is 58 grams. Take 58 grams of NaCl and add sufficient water to prepare one liter of solution.

O. Organic compounds - These are chemical compounds that contain carbon with a few exceptions (carbonates, cyanide, etc.). Carbon can form covalent bonds with a large number of atoms including itself. This makes possible long chains and rings of carbon atoms that provide the structural diversity characteristic of organic compounds.

1. Monomers and polymers - Many large organic molecules are made up of repeating subunits. These subunits are usually small organic molecules. These subunits are termed monomers and the large, chain-like, molecules which they make up are termed polymers.

a. Dehydration synthesis (condensation) - This is a special type of chemical reaction in which two monomers are bonded together. Beside the molecule which is formed, one product is always water, and therefore the name of the reaction. Virtually all biological polymers are built by dehydration synthesis.

b. Hydrolysis - This is the opposite reaction of dehydration synthesis. In this reaction a molecule of water is added across the bond holding the two monomers together and as a result the bond is broken and the monomers separate. Note that the reaction always requires and input of water.


2. Classes of organic compounds - There are four classes of organic compounds which are significant for living systems. These are carbohydrates, lipids, proteins, and nucleic acids.

a. Carbohydrates - These compounds are composed of carbon, hydrogen, and oxygen. They include the sugars, starches, glycogen, and cellulose.

(1) Carbohydrates function as fuel molecules (energy sources) and as structural parts of other molecules.

(2) Carbohydrates include both monomers and polymers. Monomers include the simple sugars, compounds containing 3 to 7 carbon atoms. Simple sugars are termed monosaccharides. Two monosaccharides can be bonded together by dehydration synthesis to form double sugars termed disaccharides. Polymers of monosaccharides are termed polysaccharides. These include starch, glycogen, and cellulose.

b. Lipids - These compounds are composed of carbon, hydrogen, and oxygen. The ratio is different from carbohydrates in that lipids contain more hydrogen per carbon.

(1) Lipids contain the fats, oils, waxes, and sterols. They are structurally heterogeneous and are grouped together based upon the fact that they will not dissolve in polar solvents but will dissolve in non-polar solvents such as acetone and benzene.

(2) Functions of the lipids in the body are as follows.

(a) Long term energy storage.

(b) Membrane components.

(c) Hormones

(3) Because of the structural diversity of lipids there is not one that can be considered typical. Neutral fats are an important class of lipids and will be used as the example. Neutral fats are sometimes referred to as triglycerides. They are composed of a molecule of glycerol plus three molecules of fatty acid, bonded together by dehydration bonds.

c. Proteins - These molecules are composed of carbon, hydrogen, oxygen, nitrogen, and usually sulfur. They are the most abundant organic molecule in the average body.

(1) Functions

(a) Structural (cell components)

(b) Contraction (muscle)

(c) Protection (immunoglobulins)

(d) Hormones

(e) Enzymes

(2) Structure - Proteins are large polymers made up of monomers known as amino acids. There are 20 different kinds of amino acids, but all have a common structure. Every amino acid contains an alpha ("first") carbon which has attached to it a carboxyl (acid) group, an amino group, a hydrogen atom, and a "R" group which is a chemical group that defines the amino acid. There are 20 "R" groups and therefore 20 different amino acids.

(3) Proteins are formed by bonding amino acids together by means of dehydration synthesis. The nitrogen of the amino group of one amino acid bonds to the carbon of the acid group on an adjacent amino acid. The C-N bond is termed the peptide bond and is characteristic of proteins.

(4) Chains of amino acids are named based upon the number of acid residues making up the chain.

(a) dipeptide - two amino acids

(b) tripeptide - three amino acids

(c) octapeptide - eight amino acids

(d) polypeptide - many amino acids

(e) protein - polypeptide with a molecular weight of at least 25,000.

(5) Proteins are very large and complex molecules. They have up to four distinct levels of structure.

(a) Primary - The kinds and sequences of amino acids.

(b) Secondary - This results from hydrogen bonding between the hydrogen of the alpha amino group to the oxygen of every other carboxyl group. This level frequently results in the twisting of the chain into a helix (corkscrew).

(c) Tertiary - This is a complex twisting and folding of the chain brought about due to hydrogen bonding and other weak bonds that form between the R groups.

(d) Quaternary - This occurs when several different polypeptide chains associate to form a single protein molecule.

    1. It is the secondary, tertiary and quaternary
    2. present) levels which give proteins their biological activity. Based upon their structure, two major groups of proteins can be recognized.

      1. Fibrous proteins - These are extended, strand like structures usually having only secondary (but some also have quaternary) structure. These proteins are highly stable and are insoluble in water. They provide mechanical strength and are the chief supporting molecules of the body. Because of their structural roles they are sometimes called structural proteins. Examples include collagen (must abundant protein in the body), keratin, and elastin.
      2. Globular proteins These are some what spherical in shape and have tertiary levels of structure and sometimes quaternary as well. These are water soluble and chemically active. They include enzymes, antibodies, hormones, and others. Because they have such a significant role in body chemistry they are sometimes termed functional proteins.
      3. Denaturization Fibrous proteins are very stable, but globular proteins are not. This is largely due to the sensitivity of the tertiary structure to disruption. The hydrogen bonds and other weak bonds that maintain the complex three dimensional structure are very easily disrupted by changes in temperature, pH, electrolyte concentration, etc. This results in shape change and loss of biological activity. This process is termed denaturizatiion. If the shape change is not to severe, restoration of initial conditions can result in the protein assuming its original shape and function. If to severe, this the protein has been irreversibly denatured and will not regain its function. An example is what happens to egg white when the egg is boiled.

(7) The sensitivity of the tertiary structure to environmental changes is the major reason that homeostasis must be maintained.

d. Nucleic acids - These are made up of carbon, oxygen, hydrogen, nitrogen, and phosphorous. They consist of long polymers made up of monomers known as nucleotide. Each nucleotide consists of a five carbon sugar (deoxyribose or ribose), a nitrogen containing base group, and a phosphate group.

(1) Functions - The nucleic acids function in the transmission of genetic information from generation to generation, and in the regulation of protein synthesis. They will be discussed in detail later.

P. Enzymes - These are biological catalysts. A catalyst is a substance which will speed up a chemical reaction but is not altered by the reaction.

1. Enzymes speed up biochemical reactions in the body. Without enzymes, no reaction could proceed at body temperature at a fast enough rate to sustain life.

2. Each and every biochemical reaction in the body has a specific enzyme which catalyzes it. The kinds of reactions which a cell can execute depends upon the kinds of enzymes which it has. The kinds of reaction which a cell can execute determines the kind of cell it will be.

3. All enzymes are proteins. Most enzymes require another small organic (non-protein) molecule in order to operate. This molecule is termed a coenzyme. Many vitamins serve as coenzymes.

4. Every enzyme has an active site. This is the part of the enzyme which just fits the reactants. The active site is due to the tertiary and structure of the enzyme. Anything that alters this level will alter enzyme activity and the altered activity will result in a slowdown or cessation of the chemical reaction being catalyze. This is why it is essential to maintain constant conditions (homeostasis) for the cells of the body.

5. Enzymes function by lowering the energy of activation of a reaction.