Chemistry of Life

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Chemistry of Life

WHY YOU NEED TO KNOW

HISTORY

Before Aureolus Paracelsus (Philippus Theophrastus Bombastus von Hohenheim; 1493–1541 ce) the principles of Western medical practice evolved primarily from witchcraft, folk remedies, and religious mysticism. An organized step forward had occurred earlier with the Greek physician Hippocrates (460–377 bce), known as the father of medicine, and Galen (131–201 ce), another Greek physician who practiced during the Roman era. Hippocrates accepted some of the rational concepts of his predecessors and taught that the body has natural resources to respond to disease and injury. Moreover, he believed that recovery from injury or disease is best implemented by getting the body in a condition to heal itself by using fresh air, water, and healthy food from nature, with limited intervention in the form of massages, purges, enemas, therapies, or drugs. His concepts of medical practice were based on a balance of the so-called “four humors”: blood, phlegm, black bile, and yellow bile, to which Galen added the four “elements”: earth, air, fire, and water. In addition, Galen introduced numerous medicaments called “galenicals,” some of which are still in use today (i.e., Galen’s cerate or cold cream) and an alcoholic extract called tincture of opium to alleviate pain.

Paracelsus in a much later era did not ascribe to these theories. Although he had studied medicine he didn’t obtain a degree to practice and was more interested in chemistry and alchemy. He proposed that the body was made up of chemicals and that disease was an imbalance of these chemicals that could be treated and/or corrected by the use of chemicals. He favored simple chemicals rather than complex compound chemical mixtures. He also prepared alcohol extracts or tinctures. Furthermore, he understood that the dose of a chemical was an important factor in determining effects that ranged from therapeutic to lethal, writing “All things are poisons, for there is nothing without poisonous qualities. It is only the dose which makes a thing a poison.” Successful, effective modern medical therapeutics stem from this chemical concept.

IMPACT

The realization and acceptance of the role of body chemistry in health and disease have impacted and shaped the understanding of current medical practice and of rational drug development. This understanding extends to the body’s native physiological responses and to its responses to drugs. For example, according to the concept of chemical molecular structure (CMS), the configuration of specialized molecules on some cells complements or recognizes and fits the configuration of certain molecules (receptors) on other cells in a lock-and-key fashion. When this chemical recognition coupling occurs it may initiate or interfere with a cascade of events that lead to a particular response. This chemical communication, modified by the genetic chemical directions given our cells, is the foundation for responses to our individual internal and external environments. If drugs such as antibiotics are administered, their degree of effectiveness is determined by how well the CMS of the antibiotic complements or fits the molecules of the receptor for that antibiotic. Thus the administration of an antibiotic, its distribution via the blood vascular system to its site of action, and the response are all understood through knowledge of chemistry. Microbiology is understood by understanding its chemistry within the network of the chemistry of life.

Atoms and Ions

All cells and organisms are made up of chemicals, and understanding the basic chemical principles is essential to understanding the structure and function of all organisms.

Elements

Knowledge of the chemistry of life begins with an understanding of those chemical principles that govern the processes occurring in matter. Matter is defined as anything that occupies space and has mass. It can be in liquid, gaseous, or solid form and is composed of elements, the smallest particles of which are atoms. Elements cannot be broken down further by natural forces. Oxygen, carbon, hydrogen, nitrogen, phosphorus, and sulfur are some of the elements most commonly found in living cells (Table 2.1). Although these chemical elements usually do not exist in free form, they do occur in combinations called chemical compounds. The shorthand expression of a chemical compound is its chemical formula. For example, the chemical formula of table salt or sodium chloride is NaCl (see Chemical Notations, below).

TABLE 2.1

Common Elements in Living Organisms

Element Symbol Atomic Number Atomic Weight
Hydrogen H 1 1
Carbon C 6 12
Nitrogen N 7 14
Oxygen O 8 16
Sodium Na 11 23
Magnesium Mg 12 24.3
Phosphorus P 15 31
Sulfur S 16 32.1
Chlorine Cl 17 35.5
Potassium K 19 39.1
Calcium Ca 20 40.1
Iron Fe 26 55.8
Cobalt Co 27 58.9
Copper Cu 29 63.5
Zinc Zn 30 65.4
Iodine I 53 126.9

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Atomic Model

All atoms have the same fundamental structure consisting of a center, or atomic nucleus, and surrounding shells (Figure 2.1), but because of the different numbers of subatomic particles, each element has its own characteristic atomic structure. Located in the center of the atom is the atomic nucleus, which consists of positively charged particles called protons, and particles without charge called neutrons. The atomic weight (atomic mass) of an atom is equal to the sum of protons and neutrons. The atomic number indicates the number of protons in the atomic nucleus. Surrounding the atomic nucleus in shells are negatively charged subatomic particles called electrons. Electrons travel around the nucleus at high speed and occupy positions in a volume of space called an orbital or electron cloud. These orbitals form an energy level also referred to as shells, in which the electrons usually remain. Electrons fill the orbitals and shells in pairs, and each orbital within a shell can carry two electrons.

The nucleus of a given atom is surrounded by successive shells spaced further and further away from the nucleus. The energy level of electrons increases with the distance of their shells from the nucleus. The innermost (first) shell can be occupied by up to two electrons within one orbital, the second shell with up to eight within four orbitals, and each consecutive shell can potentially hold more electrons. However, most elements with biological significance need eight electrons to fill the outermost shell. The shells always fill sequentially from the inside out: two electrons in the first shell, eight in the next, and so on. For example, carbon with 6 electrons carries 2 electrons in the first shell and 4 in the second shell, and sodium with 11 electrons has 2 electrons in the first shell, 8 in the second, and 1 in the third shell (Figure 2.2).

In general, the number of protons and electrons of an atom are equal, making the atom an electrically neutral unit. The stability of an atom depends on the number of electrons in the outermost shell. For example, an atom is most stable if the outermost shell is filled to its capacity. Hydrogen is the simplest element, with the atomic number of 1, and therefore has one electron in the outermost shell. Helium, with the atomic number of 2, has two electrons in the outermost shell. This shell is fully occupied and is stable. Helium atoms will not react with each other and also cannot combine with atoms of other elements. Helium is therefore called an inert gas.

If the outermost shell is not complete, the atom can participate in a chemical reaction and form a chemical bond. Electrons in the outermost shell of an atom that are available for chemical bonding are called valence electrons. These electrons determine what kind of chemical bonds, if any, the atom can form.

Isotopes are atoms with the same number of protons but a different number of neutrons. The atomic number of isotopes is unchanged because the number of protons remains the same and only the atomic weight is different. For example, the element hydrogen has two isotopes (Figure 2.3):

Radioisotopes are unstable because of their imbalance of energy within the nucleus. When the nucleus loses a neutron it gives off energy and is said to be radioactive. Radioactivity is the release of energy and matter that results from changes in the nucleus of an atom. Tritium is an example of a radioactive isotope that is used in research and clinical procedures.

Ions

Ions are electrically charged atoms, molecules, or subatomic particles that are formed when one or more valence electrons are transferred from one atom to another (see Formation and Classification of Chemical Bonds and Forces, below). If an atom loses one or more electrons to another atom, it becomes positive (+), whereas the atom that gains the electron becomes negative (−). Positively charged ions are called cations, and in an electric field move toward the negative pole, the cathode. Negatively charged ions, referred to as anions, move toward the positive pole, or anode, of an electric field.

A substance that dissociates into free ions when dissolved in a solvent such as water is called an electrolyte. The solvent in which it is dissolved can then conduct an electric current and is referred to as an electrically conductive medium. Because these solvents contain ions or electrolytes they are called ionic solutions. Chemically they are acids, bases, or salts (see Acids, Bases, and the pH Scale, below).

Several cations and anions (Table 2.2) are important components of higher life forms. All of these higher life forms require a complex electrolyte balance, called an osmotic gradient, between their intercellular and extracellular fluid compartments (see Chapter 3, Cell Structure and Function). This maintenance of a precise internal balance of electrolytes to maintain the osmotic gradient is called homeostasis. It is required to regulate the hydration, blood pH, and nerve and muscle function of an organism.

TABLE 2.2

Common Ions in Living Organisms

Cations Anions
Sodium (Na+) Chloride (Cl)
Potassium (K+ ) Bicarbonate (HCO3)
Calcium (Ca2+ ) Phosphate (PO43−)
Magnesium (Mg2+ ) Sulfate (SO42−)

Chemical Bonds and Molecules

Molecules are two or more atoms linked together by chemical bonds formed by their valence electrons. As stated above, atoms are most stable when their outermost shell is filled with eight electrons. This is the octet principle. The number of bonds a single atom can have is dependent on how many electrons are needed to complete the outermost shell. Hydrogen with one electron in the outermost shell can form one chemical bond; oxygen with eight electrons (2 + 6) (and therefore six in the outermost shell) can form two bonds; and carbon with six electrons (2 + 4) (four in the outermost shell) can form up to four chemical bonds to fill the outermost shell. When the outermost shell is not completely occupied with electrons, the atom has the tendency to interact with other atoms forming chemical bonds to achieve higher stability. These atoms then become stable and cannot react with others.

Formation and Classification of Chemical Bonds and Forces

Molecules made from atoms of different elements are called compounds. Compounds are new chemicals with properties that are different from those of the atoms of which they are composed. Groups of atoms that consistently form specific groups within compounds are referred to as functional groups. They have specific characteristics that are different from those of the individual participating atoms of that given group. Some molecules have more than one functional group, which may differ from one another. The most common functional groups found in molecules important to living organisms are shown in Table 2.3.

TABLE 2.3

Common Functional Groups in Living Organisms

Functional Group Formula Functional Group Formula
Acetyl CH3 Ethyl C2H5
Aldehyde CHO Hydroxyl OH
Amino NH2 Keto CO
Ammonium NH4 Methyl CH3
Bicarbonate HCO3 Nitrate NO3
Carbonate CO3 Phosphate PO4
Carboxyl COOH Sulfate SO4

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The principal types of chemical bonds formed by the interactions of atoms and/or molecules are covalent bonds, ionic bonds, hydrogen bonds, and those based on van der Waals forces. Covalent and hydrogen bonds occur between atoms to form a molecule, whereas hydrogen bonds and van der Waals forces are intermolecular connections. Chemical bonds vary in their strength but, in general, covalent bonds are considered the strongest bond, followed by ionic bonds, hydrogen bonds, and—with the weakest connection—the van der Waals forces.

Covalent bonds result from a sharing of electrons between two atoms of the same element or between atoms of different elements. In bonds between identical atoms such as oxygen and hydrogen, the electrons are shared equally by each atom. Covalent bonds usually are the strongest chemical bonds. Because the electrons are equally distributed, the resulting molecule is nonpolar and the bond is called a nonpolar covalent bond (Figure 2.4, A). Carbon atoms play a significant role in large organic molecules because they form stable nonpolar covalent bonds with each other. This stable framework is the backbone of organic carbon-based molecules, providing the chemical foundation of organic chemistry and of life.

The covalent bonds between atoms of two different-sized elements are polar covalent bonds, in which the electrons are unequally distributed because they are pulled toward the larger atom. As a result, one end of the molecule becomes more negative compared with the other end (Figure 2.4, B). Oxygen, nitrogen, and phosphorus atoms have a tendency to form polar covalent bonds. Polar covalent bonds are somewhat weaker than nonpolar covalent bonds. Coordinate covalent bonds, such as occurs in the formation of the ammonium ion from ammonia, are formed when both electrons are from one atom. The molecule no living organism can exist without is water, in which the atoms hydrogen and oxygen are held together by polar covalent bonds. Some properties of water result from this type of bond.

Depending on the number of electrons shared, molecules can be formed from a single covalent bond by sharing one pair of electrons, such as the bond between hydrogen atoms. Single covalent bonds are indicated by one solid line (HH). Double covalent bonds are formed by sharing two pairs of electrons, as seen between oxygen atoms. These bonds are indicated by two solid lines (OO). Triple covalent bonds may occur through the sharing of three pairs of electrons, such as between nitrogen atoms. These bonds are identified by three solid lines (NN).

Ionic bonds are formed when one or more electrons from one atom are transferred to another. If an atom loses one electron in the process it will have a charge of +1; if two electrons are lost the charge will be +2, because the protons in the nucleus will be unbalanced by the remaining electrons. The resulting anions and cations in an ionic bond are held together by attraction of their opposite charges and form an ionic compound. Ionic bonds can easily dissociate (break down) in water to form electrolyte solutions. For example, in water metals such as Na+, readily give up electrons, and nonmetals such as Cl readily take up electrons (Na+ + Cl → NaCl). If the water is evaporated, a solid crystal of NaCl, common table salt, is formed. Sodium with a total of 11 electrons (2 + 8 + 1) has only 1 electron in its outermost shell, whereas chlorine with a total of 17 electrons (2 + 8 + 7) only needs 1 electron to fill its outermost shell. The only electron in sodium’s outermost shell is therefore attracted to chlorine’s outermost shell and its transfer forms an ionic compound (Figure 2.5, A). The charged sodium chloride molecules and other salts form characteristic large crystal structures in which the atoms of the molecules alternate in a regular, geometric pattern (Figure 2.5, B). In water, NaCl readily dissociates to form an electrolyte.

Hydrogen bonds are weak chemical bonds with only about 5% of the strength of covalent bonds. However, when many hydrogen bonds are formed between two molecules, the resulting union can be strong enough to be stable. These bonds are formed by attraction forces between charged atoms within a large molecule or between adjacent molecules (Figure 2.6). Hydrogen bonds always involve a hydrogen atom with a slight positive charge and an oxygen or nitrogen atom with a slightly negative charge. Although hydrogen bonds do not form molecules they can alter the shapes of molecules or hold together different molecules. Examples of hydrogen bonds include bonds between water molecules, acetic acid molecules, amino acid molecules, and nucleic acid molecules. Hydrogen bonds are always indicated by dotted lines (—). The attraction created by hydrogen bonds keeps water in the liquid state over a wide range of temperatures.

Van der Waals forces