The Handy Chemistry Answer Book. Justin P. Lomont

solution feel the effects of the charges present. The lowest possible dielectric constant exists in a vacuum in which there is no material present to screen the charge of a field.

      

      The top graphic (A) illustrates the pi-orbital formation from two p-orbitals; the bottom graphic (B) illustrates the formation of sigma- and pi-molecular orbitals from two sp² hybridized carbon atoms.

      What is valence bond theory?

      Valence bond theory is one of two main theories (the other being molecular orbital theory) that is used to explain bonding in molecules. Valence bond theory explains bonding by describing the interactions of atomic orbitals on individual atoms as they come together to form chemical bonds. The basic idea is that orbitals with the right shapes to overlap strongly with each other will form the strongest chemical bonds. Today, valence bond theory’s description of chemical bonding based on atomic orbitals has become less popular in favor of molecular orbital theory.

      What are molecular orbitals?

      Molecular orbitals are different from atomic orbitals in that they cover several atoms and possibly even a whole molecule. While atomic orbitals originate from a single atom, molecular orbitals are formed from combinations of the atomic orbitals. Because they allow electrons to occupy the space between the atoms in a molecule, they can provide a very useful description of chemical bonds holding atoms together.

      What is molecular orbital theory?

      Molecular orbital theory is the other main theory (the first was valence bond theory) used to explain and predict bonding properties in molecules. Molecular orbital theory describes bonding interactions by using molecular orbitals that are spread out over multiple atoms, and this allows an electron’s location to be described by an orbital that bonds atoms together in a more realistic way than valence bond theory.

      What are some common structures/geometries for molecules?

      The study of chemistry has benefited greatly from knowledge of properties relating to the geometries and, especially, the symmetries of molecules. To get a sense of what shapes molecules adopt, it’s worth taking a look at a few of the geometries that come up often in the study of chemistry.

      One commonly encountered geometry is that of a tetrahedron. Methane has the molecular formula CH₄ and exists in a tetrahedral geometry with angles of approximately 109 degrees between each pair of C–H bonds.

      Linear geometries are also relatively common. Carbon dioxide has the molecular formula CO₂ and exists in a linear geometry with a 180-degree angle between the CO bonds.

      One last geometry we’ll look at here is a planar geometry. The molecule BH₃ provides one example of a planar geometry, and in this case the BH bonds are separated by angles of 120 degrees. There are also planar molecules with four bonds in a plane, and in those cases the bonds are separated by angles of 90 degrees.

      How large are molecules?

      Molecules span a wide range of sizes. The smallest molecules contain only two atoms, and these diatomic molecules have length scales that are approximately the sum of the atomic radii of the constituent atoms. The smallest molecule, H–H, has a length of only 0.74 Ångströms (7.4 × 10⁻¹¹ m). Larger molecules can be comparatively quite large. Biologically important molecules, like proteins, often contain thousands of atoms. Polymers, which are highly linked networks of covalently bonded atoms, can be even larger still, sometimes becoming so large they are visible to the naked eye.

      Is it possible to see a single molecule?

      With some of the largest single molecules, like polymers, they can actually be seen by the naked eye or through a microscope. Most molecules, however, are so small that a single isolated molecule cannot be seen with even the best microscopes. There is a physical limitation that prevents their observation with light, which has to do with the size of small molecules (lengths of ca. 0.1 to 1.0 nm) being significantly smaller than the wavelengths of visible light (400 to 700 nm). Other techniques based on diffracting electrons off of molecules, measuring the force molecules exert against a very small metal tip, and other methods have been developed to image small molecules, but it’s impossible to see most small molecules with light in the way that we conventionally see things.

      Is everything made of molecules or atoms?

      Basically, yes! The only material things that aren’t made up of atoms or molecules are the subatomic particles that make up atoms. Anything you find around your house, office, or anywhere else is made of some combination of atoms that are on the periodic table.

      How do molecules interact?

      The forces molecules exert on each other fall into a few main categories:

      Van der Waal’s interactions—Van der Waal’s interactions are the broadest group of intermolecular interactions. This includes basically all attractive and repulsive forces that don’t involve ions (charged atoms or molecules) or the rather unique situation of hydrogen bonding. Van der Waal’s interactions include forces due to the dipole moments of polar molecules as well as interactions due to induced dipoles that can form even in nonpolar molecules.

      Ionic interactions—Another class of intermolecular attractions involves attractive and repulsive forces between pairs of ions, or between ions and neutrally charged atoms or molecules. These interactions are typically stronger than those in the Van der Waal’s category. Interactions between pairs of ions are governed by Coulomb’s law, while interactions between ions and neutral molecules are either ion-dipole or ion-induced dipole interactions.

      Hydrogen bonding—A hydrogen bond is a strong interaction between a hydrogen atom and another electronegative atom (usually fluorine, oxygen, or nitrogen) that are not covalently bonded to one another. The hydrogen atom also must typically be bonded to an electronegative atom (usually oxygen or nitrogen). The origin of this strong attractive interaction is that a hydrogen atom bonded to an electronegative atom has a partial positive charge due to its lack of electron density. This allows the hydrogen atom to have a strong attractive interaction with electronegative atoms (or ions), which have partial negative charges due to their extra electron density.

      How strong are intermolecular interactions relative to a covalent bond?

      Most intermolecular interactions are fairly weak relative to a covalent bond. Covalent bonds typically involve energies on the order of 100 kilocalories per mole (a unit of energy commonly used in chemistry). Van der Waal’s interactions are the weakest type of intermolecular interaction with typical energies of roughly 0.01 to 1 kilocalories per mole (or 0.01% to 1% the strength of a covalent chemical bond) for a pair of interacting atoms. The strengths of ion-ion and ion-dipole interactions can vary widely, particularly in solutions, because the ions and/or dipoles can be separated by very different distances. The charges of ions can also be significantly shielded by solvent molecules around them. If the ions are very close together (like in a solid), their interaction energy can approach (or even exceed) that of a covalent bond. Hydrogen bonds are usually the strongest type of intermolecular interaction with energies of about 2–5 kilocalories per mole (or roughly 2% to 5% the strength of a covalent bond). Because they are such strong interactions, hydrogen bonds can play a dominant role in determining the structures of liquids, solids, and single molecules.

      What’s a solvent?

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