CoreChem:6.6 Writing Lewis Structures for Molecules
From ChemEd Collaborative
6.6 WRITING LEWIS STRUCTURES FOR MOLECULES
Because Lewis diagrams are widely used to describe the structures of molecules, it is important that you be able to construct them. The first step in this process is to draw a skeleton structure to show which atoms are linked to which. Next, determine how many valence electrons are available by adding the periodic group numbers for all atoms in the molecule. Finally, allocate the available electrons either as shared pairs or as lone pairs so that each atom has an octet (or in the case of H, a 1s2 pair).
Deciding on a Skeleton Structure
The skeleton structure of a covalent molecule can often be determined by considering the valences of the constituent atoms. Usually the atom which forms the largest number of bonds is found in the center of the skeleton, where it can connect to the maximum number of other atoms.
EXAMPLE 6.5 Hypochlorous acid has the molecular formula HOCl. Draw a structural formula.
Solution There are several possible ways to link the atoms together
The usual valence of H is 1, and so structures 3 and 4, which have two bonds to H, may be eliminated. The usual valence of Cl is also 1, and so structure 2 may also be ruled out. Structure 1 shows H forming one bond, Cl forming one, and O forming two, in agreement with the usual valences, and so it is chosen.
The total number of valence electrons available is 1 from H plus 6 from O plus 7 from Cl, or 14. Filling these into the skeleton we have
Note that O, which had the largest valence, is in the center of the skeleton.
EXAMPLE 6.6 Draw a structural formula for hydroxylamine, NH3O.
Solution In this case N has the largest valence (3), followed by O (2) and H (1). Both N and O can form “bridges” between other atoms, but H cannot. Therefore we place N and O in the center of the skeleton to give
by addition of the three H atoms.
There are a total of 5 + 3 + 6 = 14 valence electrons from N, 3H’s and O. These can be placed as follows:
Once the Lewis diagram has been determined, the molecular formula is often rewritten to remind us of what the structural formula is. For example, the molecular formula for hydroxylamine is usually written NH2OH instead of NH3O to remind us that two H’s are bonded to N and one to O. It is assumed that the person reading the formula will realize that N and O each have one valence electron left to share with each other, connecting —NH2 with —OH.
In some cases more than one skeleton structure will satisfy the valence of each atom and the octet rule as well. For example, you can verify that the molecular formula C2H6O corresponds to both of the following:
In such a case we can only decide which molecular structure we have by experiment. The properties of ethyl alcohol when diluted with water and imbibed are well known. Dimethyl ether is a gas. Like the diethyl ether used in operating rooms, it is highly explosive and can put you to sleep. Two molecules, such as dimethyl ether and ethyl alcohol, which have the same molecular formula but different structural formulas are said to be isomers.
Multiple Bonds
In order to meet the requirements of normal valence, it is sometimes necessary to have more than one bond, that is, more than one shared pair of electrons between two atoms. A case in point is formaldehyde, CH2O. In order to provide carbon with four bonds in this molecule, we must consider carbon as forming two bonds to the oxygen as well as one to each of the two hydrogens. At the same time the oxygen atom is also provided with the two bonds its normal valence requires:
Note that all four of the shared electrons in the carbon-oxygen bond are included both in the octet of carbon and in the octet of oxygen. A bond involving two electron pairs is called a double bond.
Occasionally the usual valences of the atoms in a molecule do not tell us what the skeleton structure should be. For example, in carbon monoxide, CO, it is hard to see how one carbon atom (usual valence of 4) can be matched with a single oxygen atom (usual valence of 2). In a case like this, where the valences appear to be incompatible, counting valence electrons usually leads to a structure which satisfies the octet rule. Carbon has 4 valence electrons and oxygen has 6, for a total of 10. We want to arrange these 10 electrons in two octets, but two separate groups of 8 electrons would require 16 electrons. Only by sharing 16 – 10, or 6, electrons (so that those 6 electrons are part of each octet, and, in effect, count twice) can we satisfy the octet rule. This leads to the structure
Here three pairs of electrons are shared between two atoms, and we have a triple bond.
Double and triple bonds are not merely devices for helping to fit Lewis diagrams into the octet theory. They have an objective existence, and their presence in a molecule often has a profound effect on how it reacts with other molecules. Triple bonds are invariably shorter than double bonds, which in turn are shorter than single bonds. In 
, for instance, the carbon-oxygen distance is 113 pm, in 
it is 122 pm, while in both ethyl alcohol and dimethyl ether it is 143 pm. This agrees with the wave-mechanical picture of the chemical bond as being caused by the concentration of electron density between the nuclei. The more pairs of electrons which are shared, the greater this density and the more closely the atoms are pulled together. In line with this, we would also expect multiple bonds to be stronger than single bonds. Indeed, the bond energy of C—O is found experimentally to be 360 kJ mol–1, while that of 
is 736 kJ mol–1, and that of is a gigantic 1072 kJ mol–1. The triple bond in carbon monoxide turns out to be the strongest known covalent bond.
The formation of double and triple bonds is not as widespread among the atoms of the periodic table as one might expect. At least one of the atoms involved in a multiple bond is almost always C, N, or O, and in most cases both atoms are members of this trio. Other elements complete their octets by forming additional single bonds rather than multiple bonds.
EXAMPLE 6.7 Draw structural formulas for (a) CO2 and SiO2.
Solution
a) Carbon requires four bonds, and each oxygen requires two bonds, and so two C==O double bonds will satisfy the normal valences. The structure is
b) Silicon also has a normal valence of 4, but it is not an element which readily forms double bonds. Each silicon can form single bonds to four oxygen atoms however,
Now the silicon is satisfied, but each oxygen lacks one electron and has only formed one bond. If each of the oxygens link to another silicon, they will be satisfied, but then the added silicon atoms will have unused valences:
The process of adding oxygen or silicon atoms can continue indefinitely, producing a giant lattice of covalently bonded atoms. In this giant molecule each silicon is bonded to four oxygens and each oxygen to two silicons, and so there are as many oxygen atoms as silicon. The molecular formula could be written (SiO2)n where n is a very large number. A portion of this giant molecule is shown in Fig. 6.8.
The difference in the abilities of carbon and silicon atoms to form double bonds has important consequences in the natural environment. Because C==O double bonds form readily, carbon dioxide consists of individual molecules—there are no “empty spaces” on either the carbon or oxygen atoms where additional electrons may be shared. Hence there is little to hold one carbon dioxide molecule close to another, and at ordinary temperatures the molecules move about independently. On a macroscopic scale this means that carbon dioxide has the properties of a gas. In silicon dioxide, on the other hand, strong covalent bonds link all silicon and oxygen atoms together in a three-dimensional network. At ordinary temperatures the atoms cannot vibrate far from their allotted positions, and silicon dioxide has the macroscopic properties of a solid.
As a gas, carbon dioxide is much freer than silicon dioxide to circulate through the environment. It can be removed from the atmosphere by plants in the photosynthetic process and eventually returned to the air by means of respiration. This is one of the reasons that terrestrial life is based on carbon compounds. If a supply of carbon from atmospheric carbon dioxide were not available, living organisms would be quite different in form and structure from the ones we know on earth.
Science-fiction authors are fond of suggesting, because of the periodic relationship of carbon and silicon, that life on some distant planet might be based on silicon. It is rather hard to imagine, though, the mechanism by which such life forms would obtain silicon from the rocks and soil of their planet’s surface. Certainly they would face major difficulties if the combination of silicon with oxygen to form silicon dioxide were to be used as a source of energy. Imagine breathing out a solid instead of the gaseous carbon dioxide which forms when carbon combines with oxygen during respiration in terrestrial organisms! Macroscopic properties which are determined by microscopic structure and bonding are crucial in even such fundamental activities as living and breathing.
An Excess of Bonds
There are a number of cases where the normal valences of the atoms involved do not predict the correct skeleton structure. For example, thionyl chloride, SOCl2, is found experimentally to have both chlorines and the oxygen atom bonded to sulfur:
This exceeds the usual valence of 2 for sulfur, while oxygen has one less bond than we might have expected. Molecules which deviate in this way from the usual valence rules often contain at least one atom (such as S) from the third row or below in the periodic table. One or more oxygen atoms are usually bonded just to that third-row atom instead of linking a pair of other atoms. Usually the atom which occupies the central position in the skeleton is written first in the molecular formula, although sometimes H (which forms only one bond and cannot be the central atom) precedes it. Some examples (with the central atom in italics) are: SOCl2, POCl3, HClO4, SO2Br2, and N2O. (In the last case, one of the two nitrogens occupies the central position.)
In such molecules the deviation from the normal valence occurs because at least one electron-pair bond contains two electrons which were originally associated with the same atom. Such a bond is called a coordinate covalent bond or a dative bond. An example is the bond between sulfur and oxygen in SOCl2:
Both electrons in the S―O bond were originally valence electrons of sulfur. Therefore the sharing of this electron pair adds nothing to the valence shell of sulfur, and sulfur can form one more bond than would be predicted by its normal valence. Neither electron in the coordinate covalent bond was originally associated with oxygen, and a single bond (both electrons) is sufficient to provide an octet when added to oxygen’s six valence electrons. Hence oxygen forms one less bond than expected.
It should be noted that there is nothing to distinguish a coordinate covalent bond from any other covalent bond once a structural formula has been drawn. A pair of electrons is still a pair of electrons no matter where it came from. The distinction is merely one we make when trying to fit electron pairs into octets around each atom. Structural formulas for the other examples of unusual valence we have mentioned are shown below with coordinate covalent bonds indicated in color:
It is good practice to draw out the complete Lewis diagram for each of these molecules, differentiating electrons from different nuclei with different symbols such as × and ●, and satisfy yourself that they obey the octet rule.
Polyatomic Ions
Our discussion of ionic compounds in Sec. 6.3 was confined to monatomic ions. However, more complex ions, containing several atoms but still having a positive or negative charge, occur quite frequently in chemistry. Well-known examples of such polyatomic ions are the sulfate ion (SO42–), the hydroxide ion (OH–), the hydronium ion (H3O+), and the ammonium ion (NH4+). The atoms in these ions are joined together by covalent electron-pair bonds, and we can draw Lewis structures for the ions just as we can for molecules. The only difference is that the number of electrons in the ion does not exactly balance the sum of the nuclear charges. Either there are too many electrons, in which case we have an anion, or too few, in which case we have a cation.
Consider, for example, the hydroxide ion (OH–) for which the Lewis structure is
A neutral molecule containing one O and one H atom would contain only seven electrons, six from O and one from H. The hydroxide ion, though, contains an octet of electrons, one more than the neutral molecule. The hydroxide ion must thus carry a single negative charge. In order to draw the Lewis structure for a given ion, we must first determine how many valence electrons are involved. Suppose the structure of H3O+ is required. The total number of electrons is obtained by adding the valence electrons for each atom, 6 + 1 + 1 + 1 = 9 electrons. We must now subtract 1 electron since the species under consideration is not H3O but H3O+. The total number of electrons is thus 9 – 1 = 8. Since this is an octet of electrons, we can place them all around the O atom. The final structure then follows very easily:
In more complicated cases it is often useful to calculate the number of shared electron pairs before drawing a Lewis structure. This is particularly true when the ion in question is an oxyanion (i.e., a central atom is surrounded by several O atoms). A well-known oxyanion is the carbonate ion, which has the formula CO32–. (Note that the central atom C is written first, as was done earlier for molecules.) The total number of valence electrons available in CO32– is
4(for C) + 3 × 6(for O) + 2(for the –2 charge) = 24
We must distribute these electrons over 4 atoms, giving each an octet, a requirement of 4 × 8 = 32 electrons. This means that 32 – 24 = 8 electrons need to he counted twice for octet purposes; i.e., 8 electrons are shared. The a ion thus contains four electron-pair bonds. Presumably the C atom is double-bonded to one of the O’s and singly bonded to the other two:
In this diagram the 4C electrons have been represented by dots, the 18 O electrons by ×’s, and the 2 extra electrons by colored dots, for purposes of easy reference. Real electrons do not carry labels like this; they are all the same.
There is a serious objection to the Lewis structure just drawn. How do the electrons know which oxygen atom to single out and form a double bond with, since there is otherwise nothing to differentiate the oxygens? The answer is that they do not. To explain the bonding in the CO32– ion and some other molecules requires an extension of the Lewis theory. We will pursue this matter further in the next chapter. We end this section with two examples.
EXAMPLE 6.8 Draw a Lewis structure for the molecule ethylene, C2H4.
Solution Since hydrogen atoms are univalent, they must certainly all be bonded to carbon atoms, presumably two to each carbon. Each carbon atom thus has the situation
in which two bonds must still be accounted for. By assuming that the two carbon atoms are joined by a double bond, all the valence requirements are satisfied, and we can draw a Lewis structure containing satisfactory octets:
EXAMPLE 6.9 Draw a Lewis structure for the sulfite ion, SO32–.
Solution The safest method here is to count electrons. The total number of valence electrons available is
6(for S) + 3 × 6(for O) + 2(for the charge) = 26
To make four octets for the four atoms would require 32 electrons, and so the difference, 32 – 26 = 6, gives the number of shared electrons. There are thus only three electron-pair bonds in the ion. The central S atom must be linked by a single bond to each O atom.
Note that each of the S—O bonds is coordinate covalent.
