methane, ammonia, water, and hydrogen juoride
The elements in the first main row of the periodic table are
Lithium and beryllium are able to form positive ions by loss of one or two electrons, respectively. Boron is in an intermediate position and its somewhat unusual bonding properties are considered later in the book (Section 19.5). Carbon, nitrogen, oxygen, and fluorine all have the ability to form covalent bonds because each can complete its octet by sharing electrons with other atoms. (Fluorine or oxygen can also exist as stable anions in compounds such as Na@F or Na@OH.)
The degree of sharing of electrons in a covalent bond will not be exactly equal if the elements being linked are different. The relative attractive power exerted by an element on the electrons in a covalent bond can be expressed by its electronegativity. In one quantitative definition of electronegativity we have an increase in electronegativity along the series toward fluorine as follows :
The electronegativity of hydrogen is 2.0, close to that for carbon. Each covalent bond between elements with different electronegativities will have the bonding electrons unequally shared between them, which leads to what is called polar character. In a carbon-fluorine bond the pair of electrons are attracted more to the fluorine nucleus than to the carbon nucleus. The regions of space occupied by electrons are called orbitals and, in a molecule such as CF,, the pair of electrons in the orbital that represents each covalent bond will not be divided equally between the carbon and fluorine but will be polarized towards fluorine.
We say that such a bond is dipolar and it can be represented, when necessary, by the symbols
The electronegativity of oxygen is less than that of fluorine and closer to that of carbon; therefore the polarity of a C-0 bond will be less than that of a C-F bond. Clearly the polarity of a C-N bond will be smaller still.
Even though a n~olecule contains polar bonds, the molecule itself may be nonpolar, that is, not possess a dipole moment. This will occur when the molecule has a shape (or symmetry) such that the dipoles of the individual bonds cancel each other. Thus, molecules such as F-0-F and H-0-H will have dipole moments (as they do) if the angle between the two F-0 (or H-0) bonds is different from 180°, and zero dipole moment if the angle is 180". To predict whether a molecule has a dipole moment, it is therefore
necessary to know its shape, and in the next section the principles governing the shapes of covalently bound molecules are considered, with special reference to the series CH4,, NH3,, H2,O, and HF.
If a substance is a liquid, it is an easy matter to show experimentally whether its molecules are polar and have a dipole moment. All you have to do is to hold an object carrying an electrostatic charge near a fine stream of the falling liquid and note whether the stream is deflected. The charged object can be as simple as a glass rod rubbed on silk or an amber rod rubbed on cat's fur, the charge being positive in the first case and negative in the second. A fine stream of water is sharply deflected by such an object and this shows that the individual molecules in the liquid have positive and negative ends. The molecules tend to orient themselves so that the appropriately charged end is directed towards the charged object (for example, the negative end, oxygen, toward the positively charged rod) and then the electrostatic attraction draws the molecules toward the rod. A fine stream of carbon tetra-chloride (tetrachloromethane), CCI,, cannot be deflected at all. This shows that the CCI, molecule is sufficiently symmetrical in its arrangement of the four carbon-chlorine bonds so that the polarities of these bonds cancel each other.
A. MOLECULAR SHAPES
It is important to recognize that an understanding of the shapes of organic compounds is absolutely vital to understanding the physical, chemical, and biochemical properties of organic compounds. We present in this section a few simple concepts which will turn out later to be of great utility in predicting and correlating the shapes of complex organic molecules.
The compounds CH,, NH, , H,O, and HF are all isoelectronic: they have the same number of electrons, 10. Two are in the inner K shell of the central atom and eight are in the valence, or bonding, shell. The bonding arrangements can be indicated by Lewis structures :
Carbon, nitrogen, oxygen, and fluorine have, respectively, contributed four, five, six, and seven of the electrons that make up the octet. Because no more than two electrons can occupy an orbital, we will expect that the electrons in the octet can be treated as four distinct pairs. The electron pairs repel one another and, if the four pairs are to get as far away from each other as possible, we will expect to find the four orbitals directed toward the corners of a tetrahedron, because this provides the maximum separation between the
electrons. Methane, CH4,, is in fact tetrahedral, as is tetrafluoromethane, CF4, . The three bonds in ammonia and the two bonds in water are directed at slightly different angles, 106.6" and 104.5", respectively. This is reasonable because the repulsions between the four pairs of electrons in each of these
molecules will not be the same. Thus, for water, two of the four orbitals contain protons and two do not. We expect somewhat greater repulsions between the nonbonding pairs than between bonding pairs, and this results in the angle of the bonding pairs being somewhat less than the tetrahedral value.
Replacement of any of the hydrogen atoms in the three molecules CH4, NH3 , and H2,O with another kind of group will alter the bond angles to some extent. Replacement of one or more such hydrogens by the methyl group (methane minus a hydrogen atom), CH3 - , gives the structures shown in Table 1.1. The methyl group is an especially important substituent group and can be conveniently represented in three ways, the last being a three-dimensional representation.
The four derivatives of methane shown in Table 1.1 are all hydrocarbonsthat is, they contain only carbon and hydrogen. Hence their physical and chemical properties will resemble those of methane itself. (The molecular shapes are not well represented by the structures in the table because each of the carbon atoms in these molecules will have a tetrahedral arrangement of bonds connected to it. The three-dimensional shapes of such hydrocarbons are considered in more detail in Chapter 3.) The three derivatives of ammonia are called amines and share many of the properties of ammonia; for example, like ammonia, they have dipole moments and are weak bases. A different situation exists with the derivatives of water; each of them is representative \ of a class of compounds The structure CH3,-OH is an alcohol while CH3,- 0-CH3, is an ether. The reason that water is considered to give two classes of compounds on methyl substitution can be traced to the great importance of the hydroxyl (OH) group in chemistry. Alcohols, like water, contain a
hydroxyl group whereas ethers do not. Hydroxyl groups have a great influence on molecular properties (see next section), and the properties of alcohols (ROH) and ethers (R-0-R) are quite different. (The symbol R is usually used in organic chemistry for an alkyl group, a connected group of atoms formed by removing a hydrogen atom from a hydrocarbon; a methyl group is one kind of R group.)
The bond angles at oxygen in the two compounds CH3,-0-H and CH3,-0-CH3, are somewhat greater than those found in water. This is expected because the CH3, group is larger than hydrogen, and interference between the CH3, groups in CH3,-0-CH3, is lessened by opening the C-0-C bond angle in the ether. A compromise angle is allowed in which the interference between the CH3, groups is reduced at the expense of moving the pairs of electrons on oxygen to less favorable arrangements with respect to one another. A common description of the overall change is "relief of steric hindrance between the CH3, groups by opening the C-0-C bond angle.
B. PHYSICAL PROPERTIES
The four compounds methane, ammonia, water, and hydrogen fluoride have the physical constants shown in Table 1.2. In each of these compounds the atoms are held together to form molecules by strong covalent bonds. The melting and boiling points are governed not by these powerful forces but rather by the weaker interactions that exist between molecules-intermolecular forces. Everything else being the same, the weaker such intermolecular interactions, the lower the temperature which will usually be required, first to break down the crystal lattice of the solid by melting, and then to separate the molecules to relatively large distances by boiling.
What is the origin of these weak, secondary forces that exist between neutral molecules? We shall consider two here: van der Waals forces and hydrogen bonding. Van der Waals forces, sometimes called London forces, depend in an important way on the numbers of electrons in a molecule. This means that, in general, the bigger the molecule the greater will be the various
possible intermolecular attractions and the higher the melting and boiling points will tend to be. Boiling points tend to increase regularly within a series of compounds as the molecular weight increases. Melting points, however, usually show much less regularity. This is because the stability of a crystal lattice depends so much on molecular symmetry, which largely determines the ability of the molecules to pack well in the lattice. Thus, the five hydrocarbons shown in Table 1.1 have the boiling and melting points shown in Table 1-3. In addition to experiencing van der Waals forces (dispersion forces), molecules containing certain groups are attracted to one another by hydrogen bonding. To be effective, hydrogen bonding requires the presence of an -OH, -N-H, or F-H group; in other words, a hydrogen atom joined to a I small electronegative atom. The covalent bonds to such hydrogen atoms are 8'3 6@ strongly polarized toward the electronegative atom, for example, R-0-H, and the partially positive hydrogen will be attracted toward the partially negative oxygen atom in a neighboring molecule. In the liquid state a number
of molecules may be linked together this way at any given time. These liaisons are not permanent because thermal energies of the molecules are sufficient to cause these bonds to break very rapidly (usually within milliseconds or less). Such bonds are continually being formed and broken and this leads to the description of such temporary aggregates in a hydrogen-bonded liquid as " flickering clusters."
By far the most important of the groups responsible for hydrogen bonding is the hydroxyl group, -OH. The strength of 0-H ..- 0 hydrogen bonds may be as much as one-tenth that of an ordinary carbon-carbon covalent bond.2 (See Section 2.4 on bond strengths.) The highest boiling point in the
series CH4,, NH3, , H2,O, HF belongs to water and the lowest to methane, in which hydrogen bonding is completely absent (Table 1-2).
The three oxygen compounds shown in Table 1.1 have melting and boiling points as shown in Table 1.4.
The trends are exactly opposite to those expected on the basis of molecular weight alone and are the result of having two hydrogens bonded to oxygen in each molecule of water, one in the alcohol, CH3,OH, and none in the ether, CH30CH3.
The hydroxyl group also has an important influence on solubility characteristics. The alcohol CH30H is completely miscible with water because the two kinds of molecule can form hydrogen bonds to one another. On the other hand, the ether CH30CH3 is only partly soluble in water. Its oxygen atom can interact with the protons of water but it has no OH protons itself to continue the operation. Hydrocarbons have extremely low solubilities in water. Hydrocarbon molecules would tend to interfere with the hydrogen bonding between water molecules and could offer in exchange only the much weaker van der Waals forces.
The nitrogen compounds shown in Table 1-1 have boiling and melting points as shown in Table 1.5. There is not a great deal of difference between the values for the three amines. Hydrogen bonding N-H..-N is not as effective as 0-H. . -0 and the reduction in hydrogen bonding in going from
C. ACIDITY AND BASICITY
The acidity of the four compounds methane, ammonia, water, and hydrogen fluoride increases regularly as the central atom becomes more electronegative.
The ionization constant used here for water is the customary value of 10-l4 divided by the concentration of water in pure water (55 M). The symbol KHA,, denotes the equilibrium constant for dissociation of a neutral acid HA, that is, HA--HO + A@ and KHA = [H@][AO]/[HA]. The values refer to water solution whether actually measurable in water or not and the symbol HO represents the oxonium ion H30Q. .
Methane, like most other hydrocarbons, has a negligible acidity in water. Amines resemble ammonia in being very feeble acids. Alcohols are somewhat stronger and have acidities similar to that of water.
The basicities of these four compounds follow a different pattern which is not simply the reverse of that for acidity;
The symbol KB denotes the equilibrium constant for ionization of a neutral base B, that is, B + H2,OP-- BH@ + OH0' and KB = [BH@][OHO]/[B]. Base strengths are sometimes taken to be indicated by the acid strengths of the corresponding conjugate acids. When this is done the symbol KBH* should be used to denote the process being referred to; that is, KBH* = [H@][B]/[BH@] represents the acid dissociation BHQ P H@ + B. Ordinary basic ionization constants, KB, will be used in this book.
The increase in basicity from HF to H,O to NH, is readily understandable in terms of the decreasing electronegativity of the central atom along the series. Why then is the basicity of methane so low? The reason is that this molecule has no unshared pairs of electrons available for bonding to a proton as do ammonia and the other compounds with which we have compared it.
If methane is to accept a proton to form the ion CH5,@ the carbon atom must hold five hydrogen atoms with four pairs of electrons. (There is evidence that CHsQ can be generated and detected in the gas phase in a mass spectrometer. It may also be a transient intermediate in solutions of methane in the so-called " super acids." Examples of the latter are mixtures of FS03H and SbF5, ; their protonating power far exceeds that of concentrated sulfuric acid.)