A compound contains two or more atoms of different elements joined by chemical bonds. The properties of the compound depend on the arrangement of atoms in the compound and the types of bonds between them. To help in gaining this information for covalently bonded molecules, we draw structures for the molecules such as those shown in Section 7.1 for hydrogen, ammonia, formaldehyde, and nitrogen. These structures are called Lewis structures. A Lewis structure shows each atom in the molecule or ion and its relationship to the other atoms. It also shows all bonding electrons as well as those valence electrons that are nonbonding.

A. The Arrangement of Atoms
Most often the formula of a compound is written in a way that predicts the arrangement of its atoms. Each covalently bonded molecule or ion has one central atom or a chain of central atoms. The central atom or chain of atoms will be the least electronegative element in the structure. This chain is often of carbon atoms, although it may be of nitrogen or some other nonmetal. Methane, CH₄, has one central carbon atom. Butane, C₄H₁₀, has a chain of four carbon atoms. The element symbols that directly follow the central atom in the formula indicate that these atoms are bonded to the central atom(s). Thus, methane, CH₄, has four hydrogen atoms bonded to its central carbon atom. Butane, C₄H₁₀, has ten hydrogen atoms bonded to the four carbon atoms. Formaldehyde, whose formula is written CH₂O or HCHO, has two hydrogen atoms and one oxygen atom bonded to the central carbon atom. Methyl amine, CH₃NH₂, has three hydrogen atoms bonded to the carbon atom. The carbon atom is also bonded to a nitrogen atom, and two more hydrogen atoms are also bonded to the nitrogen atom.

If more than one arrangement of atoms seems possible, we choose the one with the most symmetry. In carbon dioxide, CO₂, the atoms are arranged O - C - O, a more symmetrical arrangement than C - O - O. Similarly, sulfur trioxide, SO₃ is written as

In showing the arrangement of atoms, keep in mind two things. Hydrogen rarely bonds to more than one atom. Halogens are usually bonded to only one atom. Only in polyatomic ions or molecules, such as bromate ion, BrO₃^-, or chloric acid, HClO₃, are halogens the central atom and thus bonded to more than one other atom.

Oxyacids like H₂SO₄, H₃PO₄, HClO₃, and acetic acid are a little different. The acid hydrogens, those that are written at the beginning of the formula, are not bonded to the central atom but are bonded to oxygen. Therefore, these compounds have the atomic arrangements:

B. The Number and Placement of Electrons in Lewis Structures
The Lewis structure of a molecule or ion shows the arrangement of atoms and the distribution of electrons in that molecule or ion. In Section 7.2A, we showed how to predict the arrangement of atoms; in this section we show how to predict the number and distribution of electrons. Some of these electrons will be shared; some will be unshared. The number in each category can be determined using the following steps. Notice that in these examples all the atoms follow the octet rule. We will illustrate each step of this process using a molecule of formaldehyde, CH₂O.

1. Each atom in the molecule except hydrogen is assumed to require eight electrons; hydrogen will require only two. (See Section 5.7C1 on the octet rule exceptions.) Formaldehyde, CH₂O, will require electrons to fill 20 spaces.

   	2(2)	+	8	+	8	= 20 spaces to fill
   	hydrogen		carbon		oxygen

2. Each atom contributes its valence electrons toward filling these spaces. (Valence electrons were defined in Section 5.5D.) Formaldehyde has 12 available valence electrons:

   	2(1)	+	4	+	6	= 12 valence electrons available
   	hydrogen		carbon		oxygen

3. The difference between these numbers is the number of bonding electrons shared by two atoms. Formaldehyde will contain eight shared electrons:   20 - 12 = 8 shared electrons

4. The number of unshared electrons is the difference between the number of valence electrons available and the number of shared or bonding electrons. Formaldehyde will have 12 - 8 = 4 unshared electrons. Now we draw the Lewis structure. Having determined the arrangement of atoms, we place a pair of electrons between each set of neighboring atoms. For formaldehyde, we write
   

   which uses six electrons and leaves two more electrons to be shared. These electrons will group with another pair to form a double bond. Carbon-hydrogen bonds are always single; therefore, the double bond will be between carbon and oxygen, giving the structure

   

5. The unshared electrons are then added where needed to give an octet. In the strucure for formaldehyde, we have four unshared electrons. Oxygen still does not have an octet, so the unshared electrons are added to the oxgygen, giving
   

   This structure is the complete Lewis structure of formaldehyde.

Example Draw the Lewis structure of the following a. ethylene, C2H4   b. phosphorus trichloride, PCl3     c. methyl alcohol, CH3OH Solution a. The most symmetrical arragement for ethylene, C2H4, is: The number of electron spaces is 4(2) + 2(8) = 24 hydrogen carbon The number of valence electrons is 4(1) + 2(4) = 12 hydrogen carbon The number of bonding electrons is 24 - 12 = 12 The number of unshared electrons is 12 - 12 = 0 Using a pair of electrons for each bond predicted in the most symmetrical arrangement gives: Two bonding electrons remain unused, and each carbon atom still needs two more electrons. Two electrons can fill this need by forming a double bond between the carbons, which gives the structure: All atoms in the structures as shown satisfy the octet rule; therefore no nonbonding (unshared) electrons are needed. b. The symmetrical atomic arrangement for phosphorus trichloride, PCl3 is The number of electron spaces is 8 + 3(8) = 32 phosphorus chlorine The number of valence electrons is 5 + 3(7) 26 phosphorus chlorine The number of bonding electrons is 32 - 26 = 6 The number of unshared electrons is 26 - 6 = 20 Adding six bonding electrons to the predicted arrangement gives: Adding the unshared electrons completes the octet of each atom and gives: c. Reading the formula for methyl alcohol, CH3OH, we predict the atomic arrangement to be three hydrogen atoms and the oxygen atom bonded to the carbon atom, with another hydrogen bonded to the oxygen: The number of electron spaces available is: 8 + 3(2) + 8 + 2 = 24 carbon hydrogen oxygen hydrogen The number of valence electrons is 4 + 3(1) + 6 + 1 =14 The number of bonding electrons is 24 - 14 =10 Using these bonding electrons for single bonds in the predicted arrangement gives: The number of unshared electrons is 14 - 10 = 4 The only atom still missing an octet of electrons is oxygen, which needs four more electrons. Putting the unshared electrons in the structure around oxygen, we get a complete Lewis structure of methyl alcohol.

C. Lewis Structures of Ions
To draw the Lewis structure of an ion, we follow the same steps as for drawing the Lewis structure of a molecule with one exception: In calculating the number of valence electrons available, one additional electron is added for each negative charge on the ion or one electron is subtracted for each positive charge on the ion. The entire structure is enclosed in brackets, and the charge is shown as a superscript outside the brackets.

 

Example Draw the Lewis structure of the following ions. a. hydronium ion, H3O+   b. ammonium ion, NH4+     c.chlorate ion, ClO3- Solution a. The atomic arrangement for the hydronium ion, H3O+ is: The number of spaces is 3(2) + 8 = 14 The number of valence electrons is 3(1) + 6 - = 8 hydrogen oxygen charge The number of shared electrons is 14 - 8 = 6 The number of unshared electrons is 8 - 6 = 2 The Lewis structure is: b. The atomic arrangement for the amonium ion, NH4+ is: The number of spaces is 4(2) + 8 = 16 hydrogen nitrogen The number of valence electrons is 4(1) + 5 - 1 = 8 hydrogen nitrogen charge The number of shared electrons is 16 - 8 = 8 The number of unshared electrons is 8 - 8 = 0 The Lewis structure is: c. The most symmetrical arrangement of atoms for the chlorate ion, ClO3- is: The number of spaces is 3(8) + 8 = 32 oxygen chlorine The number of valence electrons is 3(6) + 7 + 1 = 26 oxygen chlorine charge The number of shared electrons is 32 - 26 = 6 The number of nonbonding electrons is 26 - 6 = 20 The Lewis structure of this ion is:

D. Resonance (Optional)
As chemists began to work with Lewis structures, it became more and more obvious that, for a great many molecules and ions, no single Lewis structure provided a truly accurate representation. For example, a Lewis structure for the carbonate ion, CO₃^2-, shows carbon bonded to three oxygen atoms by a combination of one double bond and two single bonds. Three possible Lewis structures for CO₃^2- are shown in Figure 7.4. Each implies that one carbon oxygen bond is different from the other two. However, this difference in bonding is not the case; rather, it has been shown that all three bonds are identical.

FIGURE 7.4 Three possible Lewis structures for the carbonate ion, CO3-2

To describe molecules and ions, like the carbonate ion, for which no single Lewis structure is adequate, the theory of resonance was developed by Linus Pauling in the 1930s. According to resonance theory, many molecules and ions are best described by drawing two or more Lewis structures and considering the real molecule or ion as a hybrid (composite) of these structures. The individual Lewis structures are called contributing structures. We show that the real structure is a hybrid of the various contributing structures by connecting them with double-headed arrows as in Figure 7.5.

FIGURE 7.5 The carbonate ion can be represented as its three possible contributing structures connected with double-headed arrows to imply resonance.

Remember that the carbonate ion, or any other compound we describe in this way, has one and only one real structure. The problem is that our systems of representation are not adequate to describe the real structures of molecules and ions. The resonance method is a particularly useful way to describe the structure of these compounds, for it retains the use of Lewis structures with electron-pair bonds. We fully realize that the carbonate ion is not accurately represented by any single contributing structure (Figure 7.4).

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