By the end of this section, you will be able to do the following:
- Define matter and elements
- Describe the interrelationship between protons, neutrons, and electrons
- Compare the ways in which electrons can be donated or shared between atoms
- Explain the ways in which naturally occurring elements combine to create molecules, cells, tissues, organ systems, and organisms
At its most fundamental level, life is made up of matter. Matter is any substance that occupies space and has mass. Elements are unique forms of matter with specific chemical and physical properties that cannot break down into smaller substances by ordinary chemical reactions. There are 118 elements, but only 98 occur naturally. The remaining elements are unstable and require scientists to synthesize them in laboratories.
Each element is designated by its chemical symbol, which is a single capital letter or, when the first letter is already “taken” by another element, a combination of two letters. Some elements follow the English term for the element, such as C for carbon and Ca for calcium. Other elements’ chemical symbols derive from their Latin names. For example, the symbol for sodium is Na, referring to natrium, the Latin word for sodium.
Four elements common to all living organisms are oxygen (O), carbon (C), hydrogen (H), and nitrogen (N). These are found in greatest quantity, but some others (e.g. phosphorous) can be critically important. In the nonliving world, elements are found in different proportions, and some elements common to living organisms are relatively rare on the earth as a whole, as Table 2.1 shows. For example, the atmosphere is rich in nitrogen and oxygen but contains little carbon and hydrogen, while the earth’s crust, although it contains oxygen and a small amount of hydrogen, has little nitrogen and carbon. In spite of their differences in abundance, all elements and the chemical reactions between them obey the same chemical and physical laws regardless of whether they are a part of the living or nonliving world.
The Structure of the Atom
To understand how elements come together, we must first discuss the element’s smallest component or building block, the atom. An atom is the smallest unit of matter that retains all of the element’s chemical properties. For example, one gold atom has all of the properties of gold, like its chemical reactivity. A gold coin is simply a very large number of gold atoms molded into the shape of a coin and contains small amounts of other elements known as impurities. We cannot break down gold atoms into anything smaller while still retaining the properties of gold.
An atom is composed of two regions. The nucleus in the atom’s center contains protons and neutrons. The atom’s outermost region holds its electrons in orbit around the nucleus, as Figure 2.2 illustrates. All atoms contain protons, electrons, and neutrons, except hydrogen (H), which is made of one proton and one electron with no neutrons.
Protons and neutrons have approximately the same mass, about 1.67 × 10-24 grams. Scientists arbitrarily define this amount of mass as one atomic mass unit (amu) or one Dalton, as Table 2.2 shows. Although similar in mass, protons and neutrons differ in their electric charge. A proton is positively charged, whereas a neutron is uncharged. Therefore, the number of neutrons in an atom contributes significantly to its mass, but not to its charge. Electrons are much smaller in mass than protons, weighing only 9.11 × 10-28 grams, or about 1/1800 of an atomic mass unit. Hence, they do not contribute much to an element’s overall atomic mass. Although not significant contributors to mass, electrons do contribute greatly to the atom’s charge, as each electron has a negative charge equal to the proton’s positive charge. In uncharged, neutral atoms, the number of electrons orbiting the nucleus is equal to the number of protons inside the nucleus. In these atoms, the positive and negative charges cancel each other out, leading to an atom with no net charge.
Accounting for the sizes of protons, neutrons, and electrons, most of the atom’s volume—greater than 99 percent—is empty space. With all this empty space, one might ask why so-called solid objects do not just pass through one another. The reason they do not is that the electrons that surround all atoms are negatively charged and negative charges repel each other.
Atomic Number and Mass
Atoms of each element contain a characteristic number of protons and electrons. The number of protons determines an element’s atomic number, which scientists use to distinguish one element from another. The number of neutrons is variable, resulting in isotopes, which are different forms of the same element that vary only in the number of neutrons they possess. Together, the number of protons and neutrons determine an element’s mass number, as Figure 2.3 illustrates. Note that we disregard the small contribution of mass from electrons in calculating the mass number. We can use this approximation of mass to easily calculate how many neutrons an element has by simply subtracting the number of protons from the mass number. Since an element’s isotopes will have slightly different mass numbers, scientists also determine the atomic mass, which is the calculated mean of the mass number for its naturally occurring isotopes. Often, the resulting number contains a fraction. For example, the atomic mass of chlorine (Cl) is 35.45 because chlorine is composed of several isotopes, some (the majority) with atomic mass 35 (17 protons and 18 neutrons) and some with atomic mass 37 (17 protons and 20 neutrons).
How many neutrons do carbon-12 and carbon-13 have, respectively?
Isotopes are different forms of an element that have the same number of protons but a different number of neutrons. Some elements—such as carbon, potassium, and uranium—have naturally occurring isotopes. Carbon-12 contains six protons, six neutrons, and six electrons; therefore, it has a mass number of 12 (six protons and six neutrons). Carbon-14 contains six protons, eight neutrons, and six electrons; its atomic mass is 14 (six protons and eight neutrons). These two alternate forms of carbon are isotopes. Some isotopes may emit neutrons, protons, and electrons, and attain a more stable atomic configuration (lower level of potential energy); these are radioactive isotopes, or radioisotopes. Radioactive decay (carbon-14 decaying to eventually become nitrogen-14) describes the energy loss that occurs when an unstable atom’s nucleus undergoes this process.
Carbon is normally present in the atmosphere in the form of gaseous compounds like carbon dioxide and methane. Carbon-14 (14C) is a naturally occurring radioisotope that is created in the atmosphere from atmospheric 14N (nitrogen) by the addition of a neutron and the loss of a proton because of cosmic rays. This is a continuous process, so more 14C is always being created. As a living organism incorporates 14C initially as carbon dioxide fixed in the process of photosynthesis, the relative amount of 14C in its body is equal to the concentration of 14C in the atmosphere. When an organism dies, it is no longer ingesting 14C, so the ratio between 14C and 12C will decline as 14C decays gradually to 14N by a process called beta decay—electrons or positrons emission. This decay emits energy in a slow process.
After approximately 5,730 years, half of the starting concentration of 14C will convert back to 14N. We call the time it takes for half of the original concentration of an isotope to decay back to its more stable form its half-life. Because the half-life of 14C is long, scientists use it to date formerly living objects such as old bones or wood. Comparing the ratio of the 14C concentration in an object to the amount of 14C in the atmosphere, scientists can determine the amount of the isotope that has not yet decayed. On the basis of this amount, Figure 2.4 shows that we can calculate the age of the material, such as the pygmy mammoth, with accuracy if it is not much older than about 50,000 years. Other elements have isotopes with different half-lives. For example, 40K (potassium-40) has a half-life of 1.25 billion years, and 235U (Uranium 235) has a half-life of about 700 million years. Through the use of radiometric dating, scientists can study the age of fossils or other remains of extinct organisms to understand how organisms have evolved from earlier species.
The Periodic Table
Electron Shells and the Bohr Model
Note that there is a connection between the number of protons in an element, the atomic number that distinguishes one element from another, and the number of electrons it has. In all electrically neutral atoms, the number of electrons is the same as the number of protons. Thus, each element, at least when electrically neutral, has a characteristic number of electrons equal to its atomic number.
In 1913, Danish scientist Niels Bohr (1885–1962) developed an early model of the atom. The Bohr model shows the atom as a central nucleus containing protons and neutrons, with the electrons in circular orbitals at specific distances from the nucleus, as Figure 2.6 illustrates. These orbits form electron shells or energy levels, which are a way of visualizing the number of electrons in the outermost shells. These energy levels are designated by a number and the symbol “n.” For example, 1n represents the first energy level located closest to the nucleus.
Electrons fill orbitals in a consistent order: they first fill the orbitals closest to the nucleus, then they continue to fill orbitals of increasing energy further from the nucleus. If there are multiple orbitals of equal energy, they fill with one electron in each energy level before adding a second electron. The electrons of the outermost energy level determine the atom’s energetic stability and its tendency to form chemical bonds with other atoms to form molecules.
Under standard conditions, atoms fill the inner shells first, often resulting in a variable number of electrons in the outermost shell. The innermost shell has a maximum of two electrons, but the next two electron shells can each have a maximum of eight electrons. This is known as the octet rule, which states, with the exception of the innermost shell, that atoms are more stable energetically when they have eight electrons in their valence shell, the outermost electron shell. Figure 2.7 shows examples of some neutral atoms and their electron configurations. Notice that in Figure 2.7, helium has a complete outer electron shell, with two electrons filling its first and only shell. Similarly, neon has a complete outer 2n shell containing eight electrons. In contrast, chlorine and sodium have seven and one in their outer shells, respectively, but theoretically they would be more energetically stable if they followed the octet rule and had eight.
An atom may give, take, or share electrons with another atom to achieve a full valence shell, the most stable electron configuration. Looking at this figure, how many electrons do elements in group 1 need to lose in order to achieve a stable electron configuration? How many electrons do elements in groups 14 and 17 need to gain to achieve a stable configuration?
Understanding that the periodic table’s organization is based on the total number of protons (and electrons) helps us know how electrons distribute themselves among the shells. The periodic table is arranged in columns and rows based on the number of electrons and their location. Examine more closely some of the elements in the table’s far right column in Figure 2.5. The group 18 atoms helium (He), neon (Ne), and argon (Ar) all have filled outer electron shells, making it unnecessary for them to share electrons with other atoms to attain stability. They are highly stable as single atoms. Because they are non reactive, scientists coin them inert (or noble gases). Compare this to the group 1 elements in the left-hand column. These elements, including hydrogen (H), lithium (Li), and sodium (Na), all have one electron in their outermost shells. That means that they can achieve a stable configuration and a filled outer shell by donating or sharing one electron with another atom or a molecule such as water. Hydrogen will donate or share its electron to achieve this configuration, while lithium and sodium will donate their electron to become stable. As a result of losing a negatively charged electron, they become positively charged ions. Group 17 elements, including fluorine and chlorine, have seven electrons in their outmost shells, so they tend to fill this shell with an electron from other atoms or molecules, making them negatively charged ions. Group 14 elements, of which carbon is the most important to living systems, have four electrons in their outer shell, allowing them to make several covalent bonds (discussed below) with other atoms. Thus, the periodic table’s columns represent the potential shared state of these elements’ outer electron shells that is responsible for their similar chemical characteristics.
Although useful to explain the reactivity and chemical bonding of certain elements, the Bohr model does not accurately reflect how electrons spatially distribute themselves around the nucleus. They do not circle the nucleus like the earth orbits the sun, but we find them in electron orbitals. These relatively complex shapes result from the fact that electrons behave not just like particles, but also like waves. Mathematical equations from quantum mechanics, which scientists call wave functions, can predict within a certain level of probability where an electron might be at any given time. Scientists call the area where an electron is most likely to be found its orbital.
Recall that the Bohr model depicts an atom’s electron shell configuration. Within each electron shell are subshells, and each subshell has a specified number of orbitals containing electrons. While it is impossible to calculate exactly an electron’s location, scientists know that it is most probably located within its orbital path. The letters s, p, d, and f designate the subshells. The s subshell is spherical in shape and has one orbital. Principal shell 1n has only a single s orbital, which can hold two electrons. Principal shell 2n has one s and one p subshell, and can hold a total of eight electrons. The p subshell has three dumbbell-shaped orbitals, as Figure 2.8 illustrates. Subshells d and f have more complex shapes and contain five and seven orbitals, respectively. We do not show these in the illustration. Principal shell 3n has s, p, and d subshells and can hold 18 electrons. Principal shell 4n has s, p, d and f orbitals and can hold 32 electrons. Moving away from the nucleus, the number of electrons and orbitals in the energy levels increases. Progressing from one atom to the next in the periodic table, we can determine the electron structure by fitting an extra electron into the next available orbital.
The closest orbital to the nucleus, the 1s orbital, can hold up to two electrons. This orbital is equivalent to the Bohr model’s innermost electron shell. Scientists call it the 1s orbital because it is spherical around the nucleus. The 1s orbital is the closest orbital to the nucleus, and it is always filled first, before any other orbital fills. Hydrogen has one electron; therefore, it occupies only one spot within the 1s orbital. We designate this as 1s1, where the superscripted 1 refers to the one electron within the 1s orbital. Helium has two electrons; therefore, it can completely fill the 1s orbital with its two electrons. We designate this as 1s2, referring to the two electrons of helium in the 1s orbital. On the periodic table in Figure 2.5, hydrogen and helium are the only two elements in the first row (period). This is because they only have electrons in their first shell, the 1s orbital. Hydrogen and helium are the only two elements that have the 1s and no other electron orbitals in the electrically neutral state.
The second electron shell may contain eight electrons. This shell contains another spherical s orbital and three “dumbbell” shaped p orbitals, each of which can hold two electrons, as Figure 2.8 shows. After the 1s orbital fills, the second electron shell fills, first filling its 2s orbital and then its three p orbitals. When filling the p orbitals, each takes a single electron. Once each p orbital has an electron, it may add a second. Lithium (Li) contains three electrons that occupy the first and second shells. Two electrons fill the 1s orbital, and the third electron then occupies the 2s orbital. Its electron configuration is 1s22s1. Neon (Ne), alternatively, has a total of ten electrons: two are in its innermost 1s orbital, and eight fill its second shell (two each in the 2s and three p orbitals). Thus it is an inert gas and energetically stable as a single atom that will rarely form a chemical bond with other atoms. Larger elements have additional orbitals, comprising the third electron shell. While the concepts of electron shells and orbitals are closely related, orbitals provide a more accurate depiction of an atom’s electron configuration because the orbital model specifies the different shapes and special orientations of all the places that electrons may occupy.
Chemical Reactions and Molecules
All elements are most stable when their outermost shell is filled with electrons, according to the octet rule. This is because it is energetically favorable for atoms to be in that configuration and it makes them stable. However, since not all elements have enough electrons to fill their outermost shells, atoms form chemical bonds with other atoms, thereby obtaining the electrons they need to attain a stable electron configuration. When two or more atoms chemically bond with each other, the resultant chemical structure is a molecule. The familiar water molecule, H2O, consists of two hydrogen atoms and one oxygen atom. These bond together to form water, as Figure 2.9 illustrates. Atoms can form molecules by donating, accepting, or sharing electrons to fill their outer shells.
Chemical reactions occur when two or more atoms bond together to form molecules or when bonded atoms break apart. Scientists call the substances used in the beginning of a chemical reaction reactants (usually on the left side of a chemical equation), and we call the substances at the end of the reaction products (usually on the right side of a chemical equation). We typically draw an arrow between the reactants and products to indicate the chemical reaction’s direction. This direction is not always a “one-way street.” To create the water molecule above, the chemical equation would be:
2H + O → H2O
An example of a simple chemical reaction is breaking down hydrogen peroxide molecules, each of which consists of two hydrogen atoms bonded to two oxygen atoms (H2O2). The reactant hydrogen peroxide breaks down into water, containing one oxygen atom bound to two hydrogen atoms (H2O), and oxygen, which consists of two bonded oxygen atoms (O2). In the equation below, the reaction includes two hydrogen peroxide molecules and two water molecules. This is an example of a balanced chemical equation, wherein each element’s number of atoms is the same on each side of the equation. According to the law of conservation of matter, the number of atoms before and after a chemical reaction should be equal, such that no atoms are, under normal circumstances, created or destroyed.
2H2O2 (hydrogen peroxide) → 2H2O (water) + O2 (oxygen)
Even though all of the reactants and products of this reaction are molecules (each atom remains bonded to at least one other atom), in this reaction only hydrogen peroxide and water are representatives of compounds: they contain atoms of more than one type of element. Molecular oxygen, alternatively, as Figure 2.10 shows, consists of two doubly bonded oxygen atoms and is not classified as a compound but as a homonuclear molecule.
Some chemical reactions, such as the one above, can proceed in one direction until they expend all the reactants. The equations that describe these reactions contain a unidirectional arrow and are irreversible. Reversible reactions are those that can go in either direction. In reversible reactions, reactants turn into products, but when the product’s concentration goes beyond a certain threshold (characteristic of the particular reaction), some of these products convert back into reactants. At this point, product and reactant designations reverse. This back and forth continues until a certain relative balance between reactants and products occurs—a state called equilibrium. A chemical equation with a double-headed arrow pointing towards both the reactants and products often denotes these reversible reaction situations.
For example, in human blood, excess hydrogen ions (H+) bind to bicarbonate ions (HCO3–) forming an equilibrium state with carbonic acid (H2CO3). If we added carbonic acid to this system, some of it would convert to bicarbonate and hydrogen ions.
However, biological reactions rarely obtain equilibrium because the concentrations of the reactants or products or both are constantly changing, often with one reaction’s product a reactant for another. To return to the example of excess hydrogen ions in the blood, forming carbonic acid will be the reaction’s major direction. However, the carbonic acid can also leave the body as carbon dioxide gas (via exhalation) instead of converting back to bicarbonate ion, thus driving the reaction to the right by the law of mass action. These reactions are important for maintaining homeostasis in our blood.
HCO3− + H+ ↔ H2CO3 ↔ CO2 + H2O
Ions and Ionic Bonds
Some atoms are more stable when they gain or lose an electron (or possibly two) and form ions. This fills their outermost electron shell and makes them energetically more stable. Because the number of electrons does not equal the number of protons, each ion has a net charge. Cations are positive ions that form by losing electrons. Negative ions form by gaining electrons, which we call anions. We designate anions by their elemental name and change the ending to “-ide”, thus the anion of chlorine is chloride, and the anion of sulfur is sulfide.
Scientists refer to this movement of electrons from one element to another as electron transfer. As Figure 2.11 illustrates, sodium (Na) only has one electron in its outer electron shell. It takes less energy for sodium to donate that one electron than it does to accept seven more electrons to fill the outer shell. If sodium loses an electron, it now has 11 protons, 11 neutrons, and only 10 electrons, leaving it with an overall charge of +1. We now refer to it as a sodium ion. Chlorine (Cl) in its lowest energy state (called the ground state) has seven electrons in its outer shell. Again, it is more energy-efficient for chlorine to gain one electron than to lose seven. Therefore, it tends to gain an electron to create an ion with 17 protons, 17 neutrons, and 18 electrons, giving it a net negative (−1) charge. We now refer to it as a chloride ion. In this example, sodium will donate its one electron to empty its shell, and chlorine will accept that electron to fill its shell. Both ions now satisfy the octet rule and have complete outermost shells. Because the number of electrons is no longer equal to the number of protons, each is now an ion and has a +1 (sodium cation) or −1 (chloride anion) charge. Note that these transactions can normally only take place simultaneously: in order for a sodium atom to lose an electron, it must be in the presence of a suitable recipient like a chlorine atom.
Ionic bonds form between ions with opposite charges. For instance, positively charged sodium ions and negatively charged chloride ions bond together to make crystals of sodium chloride, or table salt, creating a crystalline molecule with zero net charge.
Physiologists refer to certain salts as electrolytes (including sodium, potassium, and calcium), ions necessary for nerve impulse conduction, muscle contractions, and water balance. Many sports drinks and dietary supplements provide these ions to replace those lost from the body via sweating during exercise.
Covalent Bonds and Other Bonds and Interactions
Another way to satisfy the octet rule is by sharing electrons between atoms to form covalent bonds. These bonds are stronger and much more common than ionic bonds in the molecules of living organisms. We commonly find covalent bonds in carbon-based organic molecules, such as our DNA and proteins. We also find covalent bonds in inorganic molecules like H2O, CO2, and O2. The bonds may share one, two, or three pairs of electrons, making single, double, and triple bonds, respectively. The more covalent bonds between two atoms, the stronger their connection. Thus, triple bonds are the strongest.
The strength of different levels of covalent bonding is one of the main reasons living organisms have a difficult time acquiring nitrogen for use in constructing their molecules, even though molecular nitrogen, N2, is the most abundant gas in the atmosphere. Molecular nitrogen consists of two nitrogen atoms triple bonded to each other and, as with all molecules, sharing these three pairs of electrons between the two nitrogen atoms allows for filling their outer electron shells, making the molecule more stable than the individual nitrogen atoms. This strong triple bond makes it difficult for living systems to break apart this nitrogen in order to use it as constituents of proteins and DNA.
Forming water molecules provides an example of covalent bonding. Covalent bonds bind the hydrogen and oxygen atoms that combine to form water molecules, as Figure 2.9 shows. The electron from the hydrogen splits its time between the hydrogen atoms’ incomplete outer shell and the oxygen atoms’ incomplete outer shell. To completely fill the oxygen’s outer shell, which has six electrons but which would be more stable with eight, two electrons (one from each hydrogen atom) are needed: hence, the well-known formula H2O. The two elements share the electrons to fill the outer shell of each, making both elements more stable.
Polar Covalent Bonds
There are two types of covalent bonds: polar and nonpolar. In a polar covalent bond, Figure 2.12 shows atoms unequally share the electrons, which are attracted more to one nucleus than the other. Because of the unequal electron distribution between the atoms of different elements, a slightly positive (δ+) or slightly negative (δ–) charge develops. This partial charge is an important property of water and accounts for many of its characteristics.
Water is a polar molecule, with the hydrogen atoms acquiring a partial positive charge and the oxygen a partial negative charge. This occurs because the oxygen atom’s nucleus is more attractive to the hydrogen atoms’ electrons than the hydrogen nucleus is to the oxygen’s electrons. Thus, oxygen has a higher electronegativity than hydrogen and the shared electrons spend more time near the oxygen nucleus than the hydrogen atoms’ nucleus, giving the oxygen and hydrogen atoms slightly negative and positive charges, respectively. Another way of stating this is that the probability of finding a shared electron near an oxygen nucleus is more likely than finding it near a hydrogen nucleus. Either way, the atom’s relative electronegativity contributes to developing partial charges whenever one element is significantly more electronegative than the other, and the charges that these polar bonds generate may then be used to form hydrogen bonds based on the attraction of opposite partial charges. (Hydrogen bonds, which we discuss in detail below, are weak bonds between slightly positively charged hydrogen atoms to slightly negatively charged atoms in other molecules.) Since macromolecules often have atoms within them that differ in electronegativity, polar bonds are often present in organic molecules.
Nonpolar Covalent Bonds
Nonpolar covalent bonds form between two atoms of the same element or between different elements that share electrons equally. For example, molecular oxygen (O2) is nonpolar because the electrons distribute equally between the two oxygen atoms.
Figure 2.12 also shows another example of a nonpolar covalent bond—methane (CH4). Carbon has four electrons in its outermost shell and needs four more to fill it. It obtains these four from four hydrogen atoms, each atom providing one, making a stable outer shell of eight electrons. Carbon and hydrogen do not have the same electronegativity but are similar; thus, nonpolar bonds form. The hydrogen atoms each need one electron for their outermost shell, which is filled when it contains two electrons. These elements share the electrons equally among the carbons and the hydrogen atoms, creating a nonpolar covalent molecule.
Hydrogen Bonds and Van Der Waals Interactions
Ionic and covalent bonds between elements require energy to break. Ionic bonds are not as strong as covalent, which determines their behavior in biological systems. However, not all bonds are ionic or covalent bonds. Weaker bonds can also form between molecules. Two weak bonds that occur frequently are hydrogen bonds and van der Waals interactions. Without these two types of bonds, life as we know it would not exist. Hydrogen bonds provide many of the critical, life-sustaining properties of water and also stabilize the structures of proteins and DNA, the building block of cells.
When polar covalent bonds containing hydrogen form, the hydrogen in that bond has a slightly positive charge because hydrogen’s electron is pulled more strongly toward the other element and away from the hydrogen. Because the hydrogen is slightly positive, it will be attracted to neighboring negative charges. When this happens, a weak interaction occurs between the hydrogen’s δ+ from one molecule and the molecule’s δ– charge on another molecule with the more electronegative atoms, usually oxygen. Scientists call this interaction a hydrogen bond. This type of bond is common and occurs regularly between water molecules. Individual hydrogen bonds are weak and easily broken; however, they occur in very large numbers in water and in organic polymers, creating a major force in combination. Hydrogen bonds are also responsible for zipping together the DNA double helix.
Like hydrogen bonds, van der Waals interactions are weak attractions or interactions between molecules. Van der Waals attractions can occur between any two or more molecules and are dependent on slight fluctuations of the electron densities, which are not always symmetrical around an atom. For these attractions to happen, the molecules need to be very close to one another. These bonds—along with ionic, covalent, and hydrogen bonds—contribute to the proteins’ three-dimensional structure in our cells that is necessary for their proper function.
Pharmaceutical chemists are responsible for developing new drugs and trying to determine the mode of action of both old and new drugs. They are involved in every step of the drug development process. We can find drugs in the natural environment or we can synthesize them in the laboratory. In many cases, chemists change potential drugs from nature chemically in the laboratory to make them safer and more effective, and sometimes synthetic versions of drugs substitute for the version we find in nature.
After a drug’s initial discovery or synthesis, the chemist then develops the drug, perhaps chemically altering it, testing it to see if it is toxic, and then designing methods for efficient large-scale production. Then, the process of approving the drug for human use begins. In the United States, the Food and Drug Administration (FDA) handles drug approval. This involves a series of large-scale experiments using human subjects to ensure the drug is not harmful and effectively treats the condition for which it is intended. This process often takes several years and requires the participation of physicians and scientists, including chemists, to complete testing and gain approval.
An example of a drug that was originally discovered in a living organism is Paclitaxel (Taxol), an anti-cancer drug used to treat breast cancer. This drug was discovered in the bark of the pacific yew tree. Another example is aspirin, originally isolated from willow tree bark. Finding drugs often means testing hundreds of samples of plants, fungi, and other forms of life to see if they contain any biologically active compounds. Sometimes, traditional medicine can give modern medicine clues as to where to find an active compound. For example, humans have used willow bark to make medicine for thousands of years, dating back to ancient Egypt. However, it was not until the late 1800s that scientists and pharmaceutical companies purified and marketed the aspirin molecule, acetylsalicylic acid, for human use.
Occasionally, drugs developed for one use have unforeseen effects that allow usage in other, unrelated ways. For example, scientists originally developed the drug minoxidil (Rogaine) to treat high blood pressure. When tested on humans, researchers noticed that individuals taking the drug would grow new hair. Eventually the pharmaceutical company marketed the drug to people with baldness to restore lost hair.
Finally, a pharmaceutical chemist may discover negative effects or even lack of effects. In the early 1960s, inventors, doctors, and even a U.S. senator claimed anti-cancer properties of a new drug, Krebiozen, and began to market and sell it aggressively. Through the process of infrared spectrometry, FDA chemist Alma Levant Hayden and her team discovered that the “miracle drug” was nothing more than a common compound called creatine. A pharmaceutical chemist’s career may involve detective work, experimentation, and drug development, all with the goal of making human beings healthier.