Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds. Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.
Let's look at an example. The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet. The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I. When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!
In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.
While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. We now know that there are two main types of chemical bonding; ionic bonding and covalent bonding.
Ionic Bonding
In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.
For example, during the reaction of sodium with chlorine: sodium (on the left) loses its one valence electron to chlorine (on the right),
resulting in
a positively charged sodium ion (left) and a negatively charged chlorine ion (right).
Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na+ and Cl- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common: Ionic bonds form between metals and nonmetals.
In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride). Ionic compounds dissolve easily in water and other polar solvents. In solution, ionic compounds easily conduct electricity. Ionic compounds tend to form crystalline solids with high melting temperatures.
This last feature, the fact that ionic compounds are solids, results from the intermolecular forces (forces between molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule becomes blurred in ionic crystals because the solid exists as one continuous system. Forces between molecules are comparable to the forces within the molecule, and ionic compounds tend to form crystal solids with high melting points as a result.
Covalent Bonding
The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.
Polar and Nonpolar Covalent Bonding There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed.
A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule. Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells.
Polar covalent bonding simulated in water
The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.
The Dipole
Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called a dipole. The water dipole is represented by the arrow in the pop-up animation (above) in which the head of the arrow points toward the electron dense (negative) end of the dipole and the cross resides near the electron poor (positive) end of the molecule.
Although born in Cleveland, Ohio, USA, on December 9, 1919, I moved to Kentucky in 1920, and lived in Lexington through my university years. After my bachelors degree at the University of Kentucky, I entered graduate school at the California Institute of Technology in 1941, at first in physics. Under the influence of Linus Pauling, I returned to chemistry in early 1942. From then until the end of 1945 I was involved in research and development related to the war. After completion of the Ph.D., I joined the faculty of the University of Minnesota in 1946, and moved to Harvard University in 1959. Harvard's recognitions include the Abbott and James Lawrence Professorship in 1971, and the George Ledlie Prize in 1971.
The early research in borane chemistry is best summarized in my book "Boron Hydrides" (W.A. Benjamin, Inc., 1963), although most of this and late work is in several scientific journals. Since about 1960, my research interests have also been concerned with the relationship between three-dimensional structures of enzymes and how they catalyze reactions or how they are regulated by allosteric transformations.
Besides memberships in various scientific societies, I have received the Bausch and Lomb honorary science award in 1937; and, from the American Chemical Society, the Award for Distinguished Service in the Advancement of Inorganic Chemistry, and the Peter Debye Award in Physical Chemistry. Local sections of this Society have given the Harrison Howe Award and Remsen Award. The University of Kentucky presented to me the Sullivan Medallion in 1941, the Distinguished Alumni Centennial Award in 1965, and an honorary Doctor of Science degree in 1963. A Doctor Honoris Causa was awarded by the University of Munich in 1976. I am a member of the National Academy of Sciences U.S.A. and of the American Academy of Arts and Sciences, and a foreign member of the Royal Netherlands Academy of Sciences and Letters.
My other activities include tennis and classical chamber music as a performing clarinetist
Interview
Interview with Professor William N. Lipscomb by Joanna Rose, science writer, December 3, 2001.
Professor Lipscomb talks about how the directions of his research abruptly changed in 1964, about how to get confidence in ideas that nobody believes in (3:30), why music is a very important part of his life (5:40) and how a chemistry set for kids made him enter science (9:55). He also gives some advice to young students (11:26).
Banquet Speech
William Lipscomb's speech at the Nobel Banquet, December 10, 1976 Your Majesties, Your Royal Highnesses, Ladies and Gentlemen,
The award this year of the Nobel Prize in chemistry for research in pure inorganic chemistry is an important event. It is a reminder that we know even now considerably less about most of the chemical elements than we know of the chemistry of carbon or the chemistry of processes underlying life itself. It is also a reminder that the inorganic area characterized by these other elements is now being incorporated into organic chemistry and biochemistry.
On this occasion in which intellectual achievement is given its highest recognition, those of us who are being honored should remind ourselves that we are tall only because we stand on the shoulders of others: those who have gone before us showing the way, those who worked with us as our colleagues, those who supported our research giving us funds, and those who have honored us, including you, most of all.
Finally, I take special recognition of the truly international nature of science. As a personal note, I take pleasure in recalling the remarkable association that I have had with many other research groups in different countries throughout the world.
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