Intermolecular forces World of Chemistry, 2000 Updated: August 29, 2013 The attractions of molecules to each other are known as intermolecular forces to distinguish them from intramolecular forces, such as covalent and ionic bonding, which act to hold atoms together within a molecule. For very large molecules such as proteins, nucleic acids, and synthetic polymers, noncovalent interactions of groups that make up the molecules have the same characteristics as intermolecular interactions even though they may actually take place between groups within the same molecule. The interactions that produce specific shapes of enzymes, referred to as protein-folding interactions, are an example of intermolecular-like interactions within individual molecules. Intermolecular forces are most significant in liquids and solids, in which molecules are close to each other. Even in liquids and solids, intermolecular forces are only strong for molecules that are near to each other. The interactions of molecules in the liquid and solid states have significant and readily observable consequences. The strength of intermolecular forces affects such properties as boiling point, miscibility, and solubility. All neighboring molecules in liquids and solids attract each other. The nature and strength of these interactions depend on the types of groups of atoms or functional groups that comprise the molecules. All molecules interact with other molecules through London dispersion forces. London dispersion forces are the attractive forces of one transient dipole--a temporary imbalance of positive and negative charge--for another. At particular instants, even atoms that are spherical on average, such as those of the noble gases, will have greater electron density on one side of the atom than another. At that instant, the atom will possess a temporary dipole with a negative charge concentration on the side of the atom with greater electron density. If this happens in the case of an argon atom in liquid argon, for example, the argon atoms next to the one with temporary dipole would feel the effect of the dipole. An atom near the negative end of the dipole would have its own electrons slightly repelled from the negative concentration of charge, developing a dipole with its positive end near the negative charge of the original atom. An argon atom on the other side of the original temporary dipole would feel its electrons attracted to the positive end of the dipole, developing a dipole with the opposite orientation. In this way, temporary dipoles are propagated through a liquid or solid. The motion of the molecules in the liquid or solid soon disrupts the pattern, but similar events take place continually. The larger the size of atoms and the more electrons they possess, the greater the probability of forming substantial transient dipole interactions. Molecules which are non-polar and non-polar functional groups of molecules experience London dispersion only with other molecules or functional groups. London dispersion forces are one form of attraction between atoms and molecules other than those due to covalent bonds, hydrogen bonds, and electrostatic forces between particles. Similar forces are Debye forces, which occur between a permanent dipole and a temporary induced dipole, and Keesom forces, which exist between two permanent dipoles. London, Debye, and Keesom forces all constitute a type of attraction known as van der Waals force.
Another type of intermolecular force is a hydrogen bond. A hydrogen bond is a force of attraction between a proton (hydrogen nucleus) on one chemical species and an electronegative atom (such as oxygen or nitrogen) on a second chemical species. For example, the hydrogen atom in one molecule of water feels a force of attraction to an oxygen atom in a second molecule of water. That force of attraction is described as a hydrogen bond between the two molecules of water. The energy of a hydrogen bond generally ranges from about 5 kj/mol to about 30 kj/mol, in comparison to the energy of a covalent bond, which can be as much as 100 kj/mol. As such, it is the strongest of the intermolecular forces. The specific geometric requirements for hydrogen bonding are profoundly important in biochemistry, notably in the operation of the genetic code. The effects of the various types of intermolecular forces can be seen in the boiling points of compounds of elements formed with hydrogen (the hydrides) in the periodic groups containing carbon, nitrogen, oxygen, and fluorine. For all the hydrides except water, ammonia, and hydrogen fluoride, the boiling points of the hydrides of each periodic group (for example, CH4, SiH4, GeH4, and SnH4) increase as the atomic number of the cation increases. This is the normal effect due to London forces--the greater the polarizability of the electron cloud, the more the condensed phase is stabilized by transient dipoles. The hydrides of the halogens (HCl, HBr, and HI) all have permanent dipoles. They boil at temperatures higher than the hydrides of the elements in the corresponding row of the preceding group of the periodic table and also show the trend of increasing boiling point as the polarizability of their electron clouds increases. The hydrides of oxygen, nitrogen, and fluorine, however, require the input of more energy to boil than would be predicted from either London forces or dipole-dipole interactions. Water, ammonia, and hydrogen fluoride all form hydrogen bonds. Each water molecule can form four hydrogen bonds with other water molecules--two involving its two hydrogen atoms, and two involving the unshared pairs of electrons on its oxygen. Thus, the effect of hydrogen bonding on its boiling point is especially great. Molecules forming hydrogen bonds are stabilized when the molecules align properly. The kinetic energy of molecules always competes against attractive forces to determine physical properties because molecular motion opposes proper alignment. More heat is required to disrupt the stable associations in a solid containing hydrogen-bonded molecules to form a liquid, in which the molecules are more mobile and not as well aligned. Similarly, compared with non-hydrogen-bonding substances, liquids in which the molecules can hydrogen bond require more heat to form a gas, in which there is very little interaction between molecules. The stronger the intermolecular forces, the more heat energy is required to cause changes of phase. Each hydrogen bond between molecules is much weaker than the covalent bonds holding the atoms together within the molecule. It may require a large amount of heat to convert liquid water into steam, but it requires significantly more energy to break a molecule of water apart into hydrogen and oxygen atoms. The additivity of intermolecular forces explains differences in properties of larger molecules. In effect, the interaction of each group of atoms of a molecule with a group of atoms of a neighboring molecule can be considered to be independent of the interactions of other groups of atoms of the molecules. The total energy required to separate two molecules is, then, the sum of all the energies of the individual forces. The more groups and the stronger each individual force, the greater the sum of
energy of forces. Among the non-polar linear alkanes, the boiling point for a molecule with many -CH2 groups such as octane, CH3(CH2) 6CH3 is higher than that of propane CH3 CH2CH3, with octane being liquid at room temperature and propane, a gas, because of the greater number of London dispersion interactions between the octane molecules. If there are ten hydrogen bonding interactions per molecule, as there would be for a large molecule such as a sugar with many -OH groups, the total energy of interaction because of hydrogen bonds is about 300 kj mol-1, which is on the order of a strong chemical bond. Such molecules are typically solids at room temperature, while molecules containing only non-polar groups of the same molecular mass may be liquid. Life on Earth may exist because of the hydrogen bond. The physical properties of water--which covers about two thirds of Earth's surface and composes a similar proportion of the human body--are in large part a result of its extensive network of hydrogen bonds. Hydrogen bonding and intermolecular forces are the basis of the genetic code and the unique structures and shapes of the nonaqueous components of life: DNA, RNA, proteins, and other biomolecules making up living systems all owe their form and function to hydrogen bonds. The double helix of DNA is held together by hydrogen bonds. These are strong but still at least four times weaker than covalent bonds. Hydrogen bonds are strong enough to hold DNA together under most situations, but are weak enough to form and break readily to enable DNA to untwine for replication. Stabilization through hydrogen bonding can also determine the ways in which molecules arrange themselves. The hydrogen bond is a directional bond, which means there is a specific architectural relationship among molecules hydrogen bonding with each other. An important illustration of the effect of the specific geometry of hydrogen-bonded molecules is the decrease in the density of water when it freezes to form ice. Most liquids increase in density as they solidify and, therefore, solid pieces of a substance typically sink when added to the liquid substance, whereas ice floats on water. For aquatic life and indeed for all life on Earth, this anomaly of water is important. As ice forms on the surface of a body of water, it eventually forms a sheet that tends to insulate the body of water below, helping to prevent the entire body of water from freezing solid. Aquatic life that would otherwise be killed by a freeze survives. Many molecules contain some polar parts, some non-polar parts, and some functional groups that can be included in hydrogen bonds. As is true for the London forces in alkanes, polar interactions are also additive. The boiling points of propane, CH3 CH2CH3 is -44 F (-42 C), while that of butane, CH3(CH2) 2CH3 is 32 F (0 C). Adding one CH2 group increases the boiling point by 76 F (42 C). When the nonpolar -CH2- unit of propane is replaced with a polar -C=O group (this gives us acetone CH3C=OCH3), the boiling point increases by 180 F (100 C), from 44 F (-42 C) to 136 (58 C). If additional groups are added the boiling point increases accordingly. intermolecular forces also affect the properties of molecular solids and some types of covalent network solids. Molecular solids are composed of small molecules held together by London forces, dipoledipole, and hydrogen-bonding interactions. Sugar is a molecular solid, composed of glucose molecules held in crystalline array by dipole-dipole interactions. Covalent network solids are composed of very large covalently bound molecules. Graphite is a covalent network. It consists of sheets of covalently bound carbon atoms held together by van der Waals forces. A diamond is a special kind of covalent network solid because it consists entirely of covalently bonded carbon atoms. Each diamond is one molecule.
Some important observable consequences of intermolecular forces are miscibility, solubility, the heat of mixing that occurs when different liquids are combined, and capillary action. Some liquids mix and some do not. Familiar examples are the miscibility of alcohol and water and the immiscibility of oil and water, which can be explained by intermolecular forces. A portion of the alcohol molecule (-OH) hydrogen bonds with water molecules and the interaction of the two molecules is energetically favorable. The pentane molecule, however, has neither hydrogen bonding nor polar groups to be attracted to water molecules. The water molecules tend to stick together and remain separate from the pentane. Because pentane is less dense than water, it floats above it. Substances that interact strongly with water, mixing well with it or dissolving in it, are called hydrophilic (water loving). Substances that interact poorly with water and tend to remain separate are called hydrophobic (water-fearing). The effect of polarity is also evident in the solubility of compounds. For example, a bent, polar molecule like sulfur dioxide, SO2, is much more soluble in polar solvents than is linear, nonpolar carbon dioxide, CO2. This effect is also additive when molecules are complex. Alcohols with very few nonpolar groups are much more soluble in water than are alcohols with many nonpolar groups. Also, compounds with many polar groups, such as the sugars and alcohols like glycerol with three -OH groups, dissolve better in water and low molecular weight alcohols than alcohols with a lower ratio of polar to nonpolar groups. This tendency is frequently referred to as "like dissolves like"; e.g., polar solvents like water dissolve polar solutes like sugar, whereas nonpolar solutes like carbon tetrachloride dissolve nonpolar solutes like oils and other hydrocarbons. The release of heat when two hydrogen-bonding solvents are mixed is an easily observed phenomenon. If the average intermolecular force in the combined system is stronger than in the two separate liquids, heat is released. There can be a noticeable temperature rise upon mixing whenever intermolecular forces between different molecules are stronger than the interactions between like molecules, whether these associations are due to London forces, dipole-dipole interactions or hydrogen bonding. It is typically easiest to observe a heat of mixing when hydrogen bonding occurs, because this is the strongest class of intermolecular interaction. When a capillary, a narrow tube, touches the surface of a liquid, fluid rises into the tube. The extent to which a liquid rises is different for different liquids. When a narrow tube is inserted into water, the water rises in the tube. This occurs because the surface of glass is quite polar. As water molecules rise along the inside surface of the capillary, they pull up other water molecules to which they have formed hydrogen bonds. The balance of gravity and the attraction of the water for the glass surface determine the height to which the water rises. Other polar liquids besides water also rise in capillaries, but some nonpolar liquids show the opposite effect, the height of the liquid inside the capillary is less than outside. The molecules of these liquids are attracted to each other more than they are to the surface of the glass. Liquid mercury shows an especially large effect because the mercury atoms are attracted to each other much more strongly than they are attracted to the glass surface. Capillary action is also responsible for absorption of liquids into paper, such as paper towels.
Source Citation "Intermolecular forces." World of Chemistry, Gale, 2000. Science in Context, link.gale group.com/apps/doc/cv2432500385/scic?u=powd59526&xid=c0d6c6ce. Access ed 11 July 2017. Gale Document Number: GALE CV2432500385