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Learning Objectives
- Draw Lewis structures depicting the bonding in simple molecules
Thus far, we have discussed the Lewis structure of atoms and ionic compounds. We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:
The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called lone pairs) and one shared pair of electrons (written between the atoms). A dash (or line) is usually used to indicate a shared pair of electrons:
In the Lewis model, a single shared pair of electrons is a single bond. Each Cl atom interacts with eight valence electrons total: the six in the lone pairs and the two in the single bond.
Lewis Structure of Molecules: https://youtu.be/xWiFCqA9Ur0
Writing Lewis Structures with the Octet Rule
For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:
For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:
- Determine the total number of valence (outer shell) electrons among all the atoms. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.
- Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).
- Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.
- Place all remaining electrons on the central atom.
- Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.
Let us determine the Lewis structures of SiH4, \(\ce{CHO2-}\), NO+, and OF2 as examples in following this procedure:
- Determine the total number of valence (outer shell) electrons in the molecule or ion.
- For a molecule, we add the number of valence electrons on each atom in the molecule:
\(\begin{align} &\phantom{+}\ce{SiH4}\\ &\phantom{+}\textrm{Si: 4 valence electrons/atom × 1 atom = 4}\\ &\underline{\textrm{+H: 1 valence electron/atom × 4 atoms = 4}}\\ &\hspace{271px}\textrm{= 8 valence electrons} \end{align}\)
- For a negative ion, such as \(\ce{CHO2-}\), we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each single negative charge):
\(\ce{CHO2-}\\
\textrm{C: 4 valence electrons/atom × 1 atom} \hspace{6px}= \phantom{1}4\\
\textrm{H: 1 valence electron/atom × 1 atom} \hspace{12px}= \phantom{1}1\\
\textrm{O: 6 valence electrons/atom × 2 atoms = 12}\\
\underline{+\hspace{100px}\textrm{1 additional electron} \hspace{9px}= \phantom{1}1}\\
\hspace{264px}\textrm{= 18 valence electrons}\) - For a positive ion, such as NO+, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons:
\(\ce{NO+}\\
\textrm{N: 5 valence electrons/atom × 1 atom} = \phantom{−}5\\
\textrm{O: 6 valence electron/atom × 1 atom}\hspace{5px} = \phantom{−}6\\
\underline{\textrm{+ −1 electron (positive charge)} \hspace{44px}= −1}\\
\hspace{260px}\textrm{= 10 valence electrons}\) - Since OF2 is a neutral molecule, we simply add the number of valence electrons:
\(\phantom{+ }\ce{OF2}\\
\phantom{+ }\textrm{O: 6 valence electrons/atom × 1 atom} \hspace{10px}= 6\\
\underline{\textrm{+ F: 7 valence electrons/atom × 2 atoms} = 14}\\
\hspace{280px}\textrm{= 20 valence electrons}\)
- For a molecule, we add the number of valence electrons on each atom in the molecule:
- Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:)
- When several arrangements of atoms are possible, as for \(\ce{CHO2-}\), we must use experimental evidence to choose the correct one. In general, the less electronegative elements are more likely to be central atoms. In \(\ce{CHO2-}\), the less electronegative carbon atom occupies the central position with the oxygen and hydrogen atoms surrounding it. Other examples include P in POCl3, S in SO2, and Cl in \(\ce{ClO4-}\). An exception is that hydrogen is almost never a central atom. As the most electronegative element, fluorine also cannot be a central atom.
- Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons.
- There are no remaining electrons on SiH4, so it is unchanged:
- Place all remaining electrons on the central atom.
- For SiH4, \(\ce{CHO2-}\), and NO+, there are no remaining electrons; we already placed all of the electrons determined in Step 1.
- For OF2, we had 16 electrons remaining in Step 3, and we placed 12, leaving 4 to be placed on the central atom:
- Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.
- SiH4: Si already has an octet, so nothing needs to be done.
- \(\ce{CHO2-}\): We have distributed the valence electrons as lone pairs on the oxygen atoms, but the carbon atom lacks an octet:
- NO+: For this ion, we added eight outer electrons, but neither atom has an octet. We cannot add any more electrons since we have already used the total that we found in Step 1, so we must move electrons to form a multiple bond:
- This still does not produce an octet, so we must move another pair, forming a triple bond:
- In OF2, each atom has an octet as drawn, so nothing changes.
Example \(\PageIndex{1}\): Writing Lewis Structures
NASA’s Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, one of Saturn’s moons. Titan also contains ethane (H3CCH3), acetylene (HCCH), and ammonia (NH3). What are the Lewis structures of these molecules?
Solution
Calculate the number of valence electrons.
- HCN: (1 × 1) + (4 × 1) + (5 × 1) = 10
- H3CCH3: (1 × 3) + (2 × 4) + (1 × 3) = 14
- HCCH: (1 × 1) + (2 × 4) + (1 × 1) = 10
- NH3: (5 × 1) + (3 × 1) = 8
Draw a skeleton and connect the atoms with single bonds. Remember that H is never a central atom:
Where needed, distribute electrons to the terminal atoms:
- HCN: six electrons placed on N
- H3CCH3: no electrons remain
- HCCH: no terminal atoms capable of accepting electrons
- NH3: no terminal atoms capable of accepting electrons
Where needed, place remaining electrons on the central atom:
- HCN: no electrons remain
- H3CCH3: no electrons remain
- HCCH: four electrons placed on carbon
- NH3: two electrons placed on nitrogen
Where needed, rearrange electrons to form multiple bonds in order to obtain an octet on each atom:
- HCN: form two more C–N bonds
- H3CCH3: all atoms have the correct number of electrons
- HCCH: form a triple bond between the two carbon atoms
- NH3: all atoms have the correct number of electrons
Exercise \(\PageIndex{1}\)
Both carbon monoxide, CO, and carbon dioxide, CO2, are products of the combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO2 has been implicated in global climate change. What are the Lewis structures of these two molecules?
- Answer
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Fullerene Chemistry
Carbon soot has been known to man since prehistoric times, but it was not until fairly recently that the molecular structure of the main component of soot was discovered. In 1996, the Nobel Prize in Chemistry was awarded to Richard Smalley, Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C60 buckminsterfullerene molecule. An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C60. This type of molecule, called a fullerene, consists of a complex network of single- and double-bonded carbon atoms arranged in such a way that each carbon atom obtains a full octet of electrons. Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors.
Summary
Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures—especially those containing second row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron-deficient molecules, and hypervalent molecules.
Glossary
- double bond
- covalent bond in which two pairs of electrons are shared between two atoms
- free radical
- molecule that contains an odd number of electrons
- hypervalent molecule
- molecule containing at least one main group element that has more than eight electrons in its valence shell
- Lewis structure
- diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion
- Lewis symbol
- symbol for an element or monatomic ion that uses a dot to represent each valence electron in the element or ion
- lone pair
- two (a pair of) valence electrons that are not used to form a covalent bond
- octet rule
- guideline that states main group atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond
- single bond
- bond in which a single pair of electrons is shared between two atoms
- triple bond
- bond in which three pairs of electrons are shared between two atoms
