11.3 Lewis Symbols and Structures

Learning Objectives

By the end of this section, you will be able to:

  • Write Lewis symbols for neutral atoms and ions
  • Draw Lewis structures depicting the bonding in simple molecules

Thus far in this chapter, we have discussed the various types of bonds that form between atoms and/or ions. In all cases, these bonds involve the sharing or transfer of valence shell electrons between atoms. In this section, we will explore the typical method for depicting valence shell electrons and chemical bonds, namely Lewis symbols and Lewis structures.

Lewis Symbols

We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons as shown in Figure 11.3a.

The symbol for calcium is written as C a. One dot is drawn to the left of the symbol and the other dot is drawn on the right side of the symbol.
Figure 11.3a: Lewis structure of calcium atom (credit: Chemistry (OpenStax), CC BY 4.0).
Lewis symbols illustrate the number of valence electrons for each element in the third period of the periodic table as shown in Figure 11.3b.  Figure 11.3c shows the Lewis symbols for the first twenty elements of the periodic table.
A table is shown that has three columns and nine rows. The header row reads “Atoms,” “Electronic Configuration,” and “Lewis Symbol.” The first column contains the words “sodium,” “magnesium,” “aluminum,” “silicon,” “phosphorus,” “sulfur,” “chlorine,” and “argon.” The second column contains the symbols and numbers “[ N e ] 3 s superscript 2,” “[ N e ] 3 s superscript 2, 3 p superscript 1,” “[ N e ] 3 s superscript 2, 3 p superscript 2,” “[ N e ] 3 s superscript 2, 3 p superscript 3,” “[ N e ] 3 s superscript 2, 3 p superscript 4,” “[ N e ] 3 s superscript 2, 3 p superscript 5,” and “[ N e ] 3 s superscript 2, 3 p superscript 6.” The third column contains Lewis structures for N a with one dot, M g with two dots, A l with three dots, Si with four dots, P with five dots, S with six dots, C l with seven dots, and A r with eight dots.
Figure 11.3b: Lewis symbols illustrate the number of valence electrons for each element in the third period of the periodic table (credit: Chemistry (OpenStax), CC BY 4.0).

 

Lewis symbols depict the number of valence electrons for each group within the periodic table for the first 20 elements. The Lewis structuresare depicted as dots representing the valence electrons within groups 1A-7A, respectively.
Figure 11.3c: Lewis symbols for elements of the periodic table.  Review the Periodic Table of the Elements in other formats in Appendix A. (credit: graphic by Revathi Mahadevan, CC BY 4.0)

 

Lewis symbols can also be used to illustrate the formation of cations from atoms, as shown here for sodium and calcium:

Two diagrams are shown. The left diagram shows a Lewis dot structure of sodium with one dot, then a right-facing arrow leading to a sodium symbol with a superscripted plus sign, a plus sign, and the letter “e” with a superscripted negative sign. The terms below this diagram read “Sodium atom” and “Sodium cation.” The right diagram shows a Lewis dot structure of calcium with two dots, then a right-facing arrow leading to a calcium symbol with a superscripted two and a plus sign, a plus sign, and the value “2e” with a superscripted negative sign. The terms below this diagram read “Calcium atom” and “Calcium cation.”
Figure 11.3d: Formation of sodium and calcium cations shown as Lewis structures (credit: Chemistry (OpenStax), CC BY 4.0).

Likewise, they can be used to show the formation of anions from atoms, as shown here for chlorine and sulfur:

Two diagrams are shown. The left diagram shows a Lewis dot structure of chlorine with seven dots and the letter “e” with a superscripted negative sign, then a right-facing arrow leading to a chlorine symbol with eight dots and a superscripted negative sign. The terms below this diagram read, “Chlorine atom,” and, “Chlorine anion.” The right diagram shows a Lewis dot structure of sulfur with six dots and the symbol “2e” with a superscripted negative sign, then a right-facing arrow leading to a sulfur symbol with eight dots and a superscripted two and negative sign. The terms below this diagram read, “Sulfur atom,” and, “Sulfur anion.”
Figure 11.3e: Formation of chloride and sulfide anions shown as Lewis structures (credit: Chemistry (OpenStax), CC BY 4.0).

Figure 11.3f demonstrates the use of Lewis symbols to show the transfer of electrons during the formation of ionic compounds.

A table is shown with four rows. The header row reads “Metal,” “Nonmetal,” and “Ionic Compound.” The second row shows the Lewis structures of a reaction. A sodium symbol with one dot, a plus sign, and a chlorine symbol with seven dots lie to the left of a right-facing arrow. To the right of the arrow a sodium symbol with a superscripted plus sign is drawn next to a chlorine symbol with eight dots surrounded by brackets with a superscripted negative sign. One of the dots on the C l atom is red. The terms “sodium atom,” “chlorine atom,” and “sodium chloride ( sodium ion and chloride ion )” are written under the reaction. The third row shows the Lewis structures of a reaction. A magnesium symbol with two red dots, a plus sign, and an oxygen symbol with six dots lie to the left of a right-facing arrow. To the right of the arrow a magnesium symbol with a superscripted two and a plus sign is drawn next to an oxygen symbol with eight dots, two of which are red, surrounded by brackets with a superscripted two a and a negative sign. The terms “magnesium atom,” “oxygen atom,” and “magnesium oxide ( magnesium ion and oxide ion )” are written under the reaction. The fourth row shows the Lewis structures of a reaction. A calcium symbol with two red dots, a plus sign, and a fluorine symbol with a coefficient of two and seven dots lie to the left of a right-facing arrow. To the right of the arrow a calcium symbol with a superscripted two and a plus sign is drawn next to a fluorine symbol with eight dots, one of which is red, surrounded by brackets with a superscripted negative sign and a subscripted two. The terms “calcium atom,” “fluorine atoms,” and “calcium fluoride ( calcium ion and two fluoride ions )” are written under the reaction.
Figure 11.3f: Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed by atoms gaining electrons. The total number of electrons does not change (credit: Chemistry (OpenStax), CC BY 4.0).

Lewis Structures

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:

A Lewis dot diagram shows a reaction. Two chlorine symbols, each surrounded by seven dots are separated by a plus sign. The dots on the first atom are all black and the dots on the second atom are all read. The phrase, “Chlorine atoms” is written below. A right-facing arrow points to two chlorine symbols, each with six dots surrounding their outer edges and a shared pair of dots in between. One of the shared dots is black and one is red. The phrase, “Chlorine molecule” is written below.
Figure 11.3g: Two chlorine atoms bonding for form chlorine diatomic molecule, Cl2, shown as Lewis structures (credit: Chemistry (OpenStax), CC BY 4.0).

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 sometimes used to indicate a shared pair of electrons:

Two Lewis structures are shown. The left-hand structure shows two H atoms connected by a single bond. The right-hand structure shows two C l atoms connected by a single bond and each surrounded by six dots.
Figure 11.3h: Lewis structure of hydrogen diatomic molecule, H2, and chlorine diatomic molecule, Cl2, (with lone pairs) (credit: Chemistry (OpenStax), CC BY 4.0).

A single shared pair of electrons is called a single bond. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.

The Octet Rule

The other halogen molecules (F2, Br2, I2, and At2) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule.

The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). For example, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl4 (carbon tetrachloride) and silicon in SiH4 (silane). Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:

Two sets of Lewis dot structures are shown. The left structures depict five C l symbols in a cross shape with eight dots around each, the word “or” and the same five C l symbols, connected by four single bonds in a cross shape. The name “Carbon tetrachloride” is written below the structure. The right hand structures show a S i symbol, surrounded by eight dots and four H symbols in a cross shape. The word “or” separates this from an S i symbol with four single bonds connecting the four H symbols in a cross shape. The name “Silane” is written below these diagrams.
Figure 11.3i: Lewis structures of carbon tetrachloride, CCl4, and silane, SiH4. Bonding pairs of electrons can be as a bash or bond between atoms (credit: Chemistry (OpenStax), CC BY 4.0).

Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH3 (ammonia). Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:

Three Lewis structures labeled, “Ammonia,” “Water,” and “Hydrogen fluoride” are shown. The left structure shows a nitrogen atom with a lone pair of electrons and single bonded to three hydrogen atoms. The middle structure shows an oxygen atom with two lone pairs of electrons and two singly-bonded hydrogen atoms. The right structure shows a hydrogen atom single bonded to a fluorine atom that has three lone pairs of electrons.
Figure 11.3j: Lewis structures of ammonia, NH3, water, H2O, and hydrogen fluoride, HF (credit: Chemistry (OpenStax), CC BY 4.0).

Double and Triple Bonds

As previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. A double bond forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH2O (formaldehyde) and between the two carbon atoms in C2H4 (ethylene):

Two pairs of Lewis structures are shown. The left pair of structures shows a carbon atom forming single bonds to two hydrogen atoms. There are four electrons between the C atom and an O atom. The O atom also has two pairs of dots. The word “or” separates this structure from the same diagram, except this time there is a double bond between the C atom and O atom. The name, “Formaldehyde” is written below these structures. A right-facing arrow leads to two more structures. The left shows two C atoms with four dots in between them and each forming single bonds to two H atoms. The word “or” lies to the left of the second structure, which is the same except that the C atoms form double bonds with one another. The name, “Ethylene” is written below these structures.
Figure 11.3k: Lewis structures of formaldehyde, CH2O, and ethylene, C2H2, showing double bonds (two shared pairs of electrons) (credit: Chemistry (OpenStax), CC BY 4.0).

A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN–):

Two pairs of Lewis structures are shown and connected by a right-facing arrow. The left pair of structures show a C atom and an O atom with six dots in between them and a lone pair on each. The word “or” and the same structure with a triple bond in between the C atom and O atom also are shown. The name “Carbon monoxide” is written below this structure. The right pair of structures show a C atom and an N atom with six dots in between them and a lone pair on each. The word “or” and the same structure with a triple bond in between the C atom and N atom also are shown. The name “Cyanide ion” is written below this structure.
Figure 11.3l: Lewis structures of carbon monoxide, CO, and cyanide ion, CN–, showing triple bonds (three shared pairs of electrons) (credit: Chemistry (OpenStax), CC BY 4.0).

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 (Figure 11.3m):

Three reactions are shown with Lewis dot diagrams. The first shows a hydrogen with one red dot, a plus sign and a bromine with seven dots, one of which is red, connected by a right-facing arrow to a hydrogen and bromine with a pair of red dots in between them. There are also three lone pairs on the bromine. The second reaction shows a hydrogen with a coefficient of two and one red dot, a plus sign, and a sulfur atom with six dots, two of which are red, connected by a right facing arrow to two hydrogen atoms and one sulfur atom. There are two red dots in between the two hydrogen atoms and the sulfur atom. Both pairs of these dots are red. The sulfur atom also has two lone pairs of dots. The third reaction shows two nitrogen atoms each with five dots, three of which are red, separated by a plus sign, and connected by a right-facing arrow to two nitrogen atoms with six red electron dots in between one another. Each nitrogen atom also has one lone pair of electrons.
Figure 11.3m: Lewis structures of the formation of hydrogen bromide, HBr, hydrogen sulfide, H2S, and nitrogen, N2 (credit: Chemistry (OpenStax), CC BY 4.0).

For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:

  1. Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.
  2. 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 centre.) Connect each atom to the central atom with a single bond (one electron pair).
  3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.
  4. Place all remaining electrons on the central atom.
  5. 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, CHO2−, NO+, and OF2 as examples in following this procedure:

  1. 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:
      [latex]\begin{array}{r r l} \text{SiH}_4 & & \\[1em] & \text{Si: 4 valence electrons/atom} \times 1 \;\text{atom} & = 4 \\[1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{H: 1 valence electron/atom} \times 4 \;\text{atoms} & = 4 \\[1em] & & = 8 \;\text{valence electrons} \end{array}[/latex]
    • For a negative ion, such as 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):
      [latex]\begin{array}{r r l} {\text{CHO}_2}^{-} & & \\[1em] & \text{C: 4 valence electrons/atom} \times 1 \;\text{atom} & = 4 \\[1em] & \text{H: 1 valence electron/atom} \times 1 \;\text{atom} & = 1 \\[1em] & \text{O: 6 valence electrons/atom} \times 2 \;\text{atoms} & = 12 \\[1em] \rule[-0.5ex]{21.5em}{0.1ex}\hspace{-21.5em} + & 1\;\text{additional electron} & = 1 \\[1em] & & = 18 \;\text{valence electrons} \end{array}[/latex]
    • 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:
      [latex]\begin{array}{r r l} \text{NO}^{+} & & \\[1em] & \text{N: 5 valence electrons/atom} \times 1 \;\text{atom} & = 5 \\[1em] & \text{O: 6 valence electrons/atom} \times 1 \;\text{atom} & = 6 \\[1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & -1 \;\text{electron (positive charge)} & = -1 \\[1em] & & = 10 \;\text{valence electrons} \end{array}[/latex]
    • Since OF2 is a neutral molecule, we simply add the number of valence electrons:
      [latex]\begin{array}{r r l} \text{OF}_{2} & & \\[1em] & \text{O: 6 valence electrons/atom} \times 1 \;\text{atom} & = 6 \\[1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{F: 7 valence electrons/atom} \times 2 \;\text{atoms} & = 14 \\[1em] & & = 20 \;\text{valence electrons} \end{array}[/latex]
  2. 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:)
    Four Lewis diagrams are shown. The first shows one silicon single boned to four hydrogen atoms. The second shows a carbon which forms a single bond with an oxygen and a hydrogen and a double bond with a second oxygen. This structure is surrounded by brackets and has a superscripted negative sign near the upper right corner. The third structure shows a nitrogen single bonded to an oxygen and surrounded by brackets with a superscripted plus sign in the upper right corner. The last structure shows two fluorine atoms single bonded to a central oxygen.
    Figure 11.3n: Skeletal structures (incomplete Lewis structures) of SiH4, CHO2−, NO+, and OF2. Step 2 of process to draw Lewis structure (credit: Chemistry (OpenStax), CC BY 4.0).

    When several arrangements of atoms are possible, as for 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 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 ClO4−. An exception is that hydrogen is almost never a central atom. As the most electronegative element, fluorine also cannot be a central atom.

  3. 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:
      Four Lewis structures are shown. The first shows one silicon single boned to four hydrogen atoms. The second shows a carbon single bonded to two oxygen atoms that each have three lone pairs and single bonded to a hydrogen. This structure is surrounded by brackets and has a superscripted negative sign near the upper right corner. The third structure shows a nitrogen single bonded to an oxygen, each with three lone pairs of electrons. This structure is surrounded by brackets with a superscripted plus sign in the upper right corner. The last structure shows two fluorine atoms, each with three lone pairs of electrons, single bonded to a central oxygen.
      Figure 11.3o: Lewis structure of SiH4. Skeletal structures (incomplete Lewis structures) of CHO2−, NO+, and OF2. Step 3 of process to draw Lewis structure (credit: Chemistry (OpenStax), CC BY 4.0).
  4. Place all remaining electrons on the central atom.
    • For SiH4, 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:
      A Lewis structure shows two fluorine atoms, each with three lone pairs of electrons, single bonded to a central oxygen which has two lone pairs of electrons.
      Figure 11.3p: Lewis structure of OF2. Step 4 of process to draw Lewis structure (credit: Chemistry (OpenStax), CC BY 4.0).
  5. 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.
    • CHO2−: We have distributed the valence electrons as lone pairs on the oxygen atoms, but the carbon atom lacks an octet:
      Two Lewis diagrams are shown with the word “gives” in between them. The left diagram, surrounded by brackets and with a superscripted negative sign, shows a carbon atom single bonded to two oxygen atoms, each with three lone pairs of electrons. The carbon atom also forms a single bond with a hydrogen atom. A curved arrow points from a lone pair on one of the oxygen atoms to the carbon atom. The right diagram, surrounded by brackets and with a superscripted negative sign, shows a carbon atom single bonded to an oxygen atom with three lone pairs of electrons, double bonded to an oxygen atom with two lone pairs of electrons, and single bonded to a hydrogen atom.
      Figure 11.3q: Lewis structure of CHO2−. Step 5 of process to draw Lewis structure (credit: Chemistry (OpenStax), CC BY 4.0).
    • NO+: For this ion, we added eight valence 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:
      Two Lewis diagrams are shown with the word “gives” in between them. The left diagram, surrounded by brackets and with a superscripted positive sign, shows a nitrogen atom single bonded to an oxygen atom, each with two lone pairs of electrons. The right diagram, surrounded by brackets and with a superscripted positive sign, shows a nitrogen atom double bonded to an oxygen atom. The nitrogen atom has two lone pairs of electrons and the oxygen atom has one.
      Figure 11.3r: Incomplete Lewis structure NO+. Step 5 of process to draw Lewis structure (credit: Chemistry (OpenStax), CC BY 4.0).
    • This still does not produce an octet, so we must move another pair, forming a triple bond:
      A Lewis structure shows a nitrogen atom with one lone pair of electrons triple bonded to an oxygen with a lone pair of electrons. The structure is surrounded by brackets and has a superscripted positive sign.
      Figure 11.3s: Lewis structure of NO+. Step 5 of process to draw Lewis structure (credit: Chemistry (OpenStax), CC BY 4.0).
    • In OF2, each atom has an octet as drawn, so nothing changes.

Example 11.3a

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

  1. Calculate the number of valence electrons.
    1. HCN: (1 × 1) + (4 × 1) + (5 × 1) = 10
    2. H3CCH3: (1 × 3) + (2 × 4) + (1 × 3) = 14
    3. HCCH: (1 × 1) + (2 × 4) + (1 × 1) = 10
    4. NH3: (5 × 1) + (3 × 1) = 8
  2. Draw a skeleton and connect the atoms with single bonds. Remember that H is never a central atom:Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom single bonded to three hydrogen atoms.
  3. Where needed, distribute electrons to the terminal atoms:Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom single bonded to three hydrogen atoms.
    1. HCN: six electrons placed on N.
    2. H3CCH3: no electrons remain.
    3. HCCH: no terminal atoms capable of accepting electrons.
    4. NH3: no terminal atoms capable of accepting electrons.
  4. Where needed, place remaining electrons on the central atom:Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms, each with a lone pair of electrons, single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom with a lone pair of electrons single bonded to three hydrogen atoms.
    1. HCN: no electrons remain.
    2. H3CCH3: no electrons remain.
    3. HCCH: four electrons placed on carbon.
    4. NH3: two electrons placed on nitrogen.
  5. Where needed, rearrange electrons to form multiple bonds in order to obtain an octet on each atom: Six Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. There are arrows from the lone pairs to the adjacent bond resulting in a Lewis structure with carbon single bonded to hydrogen and triple bonded to nitrogen (with one lone pair). The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms, each with a lone pair of electrons, single bonded to one another and each single bonded to one hydrogen atom. There are arrows from the lone pair on each carbon to the bond in between the two carbon atoms resulting in a Lewis structure with two carbon atoms triple bonded to each other and each single bonded to a hydrogen atom. The fourth structure shows a nitrogen atom with a lone pair of electrons single bonded to three hydrogen atoms.
    1. HCN: form two more C–N bonds.
    2. H3CCH3: all atoms have the correct number of electrons.
    3. HCCH: form a triple bond between the two carbon atoms.
    4. NH3: all atoms have the correct number of electrons.

    Exercise 11.3a

    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?

    Check Your Answer[1]

    A photograph of a ball and stick model of fullerene sitting on holder on a lab bench is pictured. Each carbon atom is bonded to three other carbon atoms, which results in the formation of a ball like structure.
    Figure 11.3t: This model shows the arrangement of carbon atoms in a fullerene (or ‘buckyball’) a large molecule made up of 60 carbon atoms. (credit: work by UCL Mathematical and Physical Science, CC BY 2.0)

    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 (Figure 11.3t), Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C60 buckminsterfullerene molecule (Figure 11.a in Chapter 11 Introduction). An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C60. This type of molecule, called a fullerene, shows promise in a variety of applications. 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.

    Richard Smalley (1943–2005), a professor of physics, chemistry, and astronomy at Rice University, was one of the leading advocates for fullerene chemistry. Upon his death in 2005, the US Senate honoured him as the “Father of Nanotechnology.”

    Learn more about Dr. Smalley by reading the article Richard E. Smalley – Facts [New Tab]  on the Nobel Prize website.

    Scientists in Action: Dr. Mario J. Molina

    A portrait of Dr. Mario J. Molina
    Figure 11.3u: Dr. Mario J. Molina (credit: work by http://science.in2pic.com, CC BY-SA 3.0)

    Dr. Mario J. Molina is a Mexican chemist who is best known for his role in the discovery of the impact that CFC’s (chlorofluorocarbons) have on the ozone layer. He was a co-recipient of the Nobel Prize in Chemistry in 1995.

    Part of his work included the hypothesis that CFCs in the upper atmosphere could produce a chlorine radical that would catalyze the destruction of ozone.

    Cl· + O3 → ClO· + O2

    ClO· + O· → Cl· + O2

    After data was collected that confirmed this hypothesis, steps were taken by countries across the world to discontinue the use of ozone-damaging compounds. The most significant actions were included in the Montreal Protocol [New Tab].

    Dr. Molina is currently a faculty member at UC San Diego and is involved in science outreach.

    Watch Dr. Molina address students about his inspiration to become a scientist in this Nobel Prize video on YouTube [New Tab].

    Exceptions to the Octet Rule

    Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:

    • Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.
    • Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.
    • Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.

    Odd-electron Molecules

    We call molecules that contain an odd number of electrons free radicals. Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at high temperatures.

    To draw the Lewis structure for an odd-electron molecule like NO, we follow the same five steps we would for other molecules, but with a few minor changes:

    1. Determine the total number of valence (outer shell) electrons. The sum of the valence electrons is 5 (from N) + 6 (from O) = 11. The odd number immediately tells us that we have a free radical, so we know that not every atom can have eight electrons in its valence shell.
    2. Draw a skeleton structure of the molecule. We can easily draw a skeleton with an N–O single bond:N–O
    3. Distribute the remaining electrons as lone pairs on the terminal atoms. In this case, there is no central atom, so we distribute the electrons around both atoms. We give eight electrons to the more electronegative atom in these situations; thus oxygen has the filled valence shell:
      A Lewis structure shows a nitrogen atom, with one lone pair and one lone electron double bonded to an oxygen atom with two lone pairs of electrons.
      Figure 11.3v: Lewis structure of NO where oxygen has been assigned the eight electrons as it is more electronegative compared to nitrogen (credit: Chemistry (OpenStax), CC BY 4.0).
    4. Place all remaining electrons on the central atom. Since there are no remaining electrons, this step does not apply.
    5. Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible. We know that an odd-electron molecule cannot have an octet for every atom, but we want to get each atom as close to an octet as possible. In this case, nitrogen has only five electrons around it. To move closer to an octet for nitrogen, we take one of the lone pairs from oxygen and use it to form a NO double bond. (We cannot take another lone pair of electrons on oxygen and form a triple bond because nitrogen would then have nine electrons:)
      A Lewis structure shows a nitrogen atom, with one lone pair and one lone electron double bonded to an oxygen atom with two lone pairs of electrons.
      Figure 11.3w: The Lewis structure of NO double bond where the electrons are rearranged to obtain octets where a double bond is formed so that nitrogen can move closer to an octet (credit: Chemistry (OpenStax), CC BY 4.0).

    Electron-deficient Molecules

    We will also encounter a few molecules that contain central atoms that do not have a filled valence shell. Generally, these are molecules with central atoms from groups 2 and 12, outer atoms that are hydrogen, or other atoms that do not form multiple bonds. For example, in the Lewis structures of beryllium dihydride, BeH2, and boron trifluoride, BF3, the beryllium and boron atoms each have only four and six electrons, respectively. It is possible to draw a structure with a double bond between a boron atom and a fluorine atom in BF3, satisfying the octet rule, but experimental evidence indicates the bond lengths are closer to that expected for B–F single bonds. This suggests the best Lewis structure has three B–F single bonds and an electron deficient boron. The reactivity of the compound is also consistent with an electron deficient boron. However, the B–F bonds are slightly shorter than what is actually expected for B–F single bonds, indicating that some double bond character is found in the actual molecule.

    Two Lewis structures are shown. The left shows a beryllium atom single bonded to two hydrogen atoms. The right shows a boron atom single bonded to three fluorine atoms, each with three lone pairs of electrons.
    Figure 11.3x: The Lewis structures for beryllium dihydride, BeH2 (left), and boron trifluoride, BF3 (right) (credit: Chemistry (OpenStax), CC BY 4.0).

    An atom like the boron atom in BF3, which does not have eight electrons, is very reactive. It readily combines with a molecule containing an atom with a lone pair of electrons. For example, NH3 reacts with BF3 because the lone pair on nitrogen can be shared with the boron atom (Figure 11.3y).

    A reaction is shown with three Lewis diagrams. The left diagram shows a boron atom single bonded to three fluorine atoms, each with three lone pairs of electrons. There is a plus sign. The next structure shows a nitrogen atom with one lone pair of electrons single bonded to three hydrogen atoms. A right-facing arrow leads to the final Lewis structure that shows a boron atom single bonded to a nitrogen atom and single bonded to three fluorine atoms, each with three lone pairs of electrons. The nitrogen atom is also single bonded to three hydrogen atoms. The bond between the boron atom and the nitrogen atom is colored red.
    Figure 11.3y: The reaction between BF3 and NH3 come together to form a single bond between boron and nitrogen because boron requires the lone pair of electrons from nitrogen to fulfill its octet (credit: Chemistry (OpenStax), CC BY 4.0).

    Hypervalent Molecules

    Elements in the second period of the periodic table (n = 2) can accommodate only eight electrons in their valence shell orbitals because they have only four valence orbitals (one 2s and three 2p orbitals). Elements in the third and higher periods (n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules. Figure 11.3z shows the Lewis structures for two hypervalent molecules, PCl5 and SF6.

    A reaction is shown with three Lewis diagrams. The left diagram shows a boron atom single bonded to three fluorine atoms, each with three lone pairs of electrons. There is a plus sign. The next structure shows a nitrogen atom with one lone pair of electrons single bonded to three hydrogen atoms. A right-facing arrow leads to the final Lewis structure that shows a boron atom single bonded to a nitrogen atom and single bonded to three fluorine atoms, each with three lone pairs of electrons. The nitrogen atom is also single bonded to three hydrogen atoms. The bond between the boron atom and the nitrogen atom is colored red.
    Figure 11.3z: The Lewis structures for two hypervalent molecules, PCl5 and SF6 (credit: Chemistry (OpenStax), CC BY 4.0).

    In some hypervalent molecules, such as IF5 and XeF4, some of the electrons in the outer shell of the central atom are lone pairs (Figure 11.3aa).

    Two Lewis structures are shown. The left shows an iodine atom with one lone pair single bonded to five fluorine atoms, each with three lone pairs of electrons. The right diagram shows a xenon atom with two lone pairs of electrons single bonded to four fluorine atoms, each with three lone pairs of electrons.
    Figure 11.3aa: The Lewis structures for hypervalent molecules, IF5 and XeF4 (credit: Chemistry (OpenStax), CC BY 4.0).

    When we write the Lewis structures for these molecules, we find that we have electrons left over after filling the valence shells of the outer atoms with eight electrons. These additional electrons must be assigned to the central atom.

    Example 11.3b

    Writing Lewis Structures: Octet Rule Violations

    Xenon is a noble gas, but it forms a number of stable compounds. We examined XeF4 earlier. What are the Lewis structures of XeF2 and XeF6?

    Solution

    We can draw the Lewis structure of any covalent molecule by following the six steps discussed earlier. In this case, we can condense the last few steps, since not all of them apply.

    1. Calculate the number of valence electrons:XeF2: 8 + (2 × 7) = 22XeF6: 8 + (6 × 7) = 50
    2. Draw a skeleton joining the atoms by single bonds. Xenon will be the central atom because fluorine cannot be a central atom:
      Two Lewis diagrams are shown. The left depicts a xenon atom single bonded to two fluorine atoms. The right shows a xenon atom single bonded to six fluorine atoms.
    3. Distribute the remaining electrons.
      1. XeF2: We place three lone pairs of electrons around each F atom, accounting for 12 electrons and giving each F atom 8 electrons. Thus, six electrons (three lone pairs) remain. These lone pairs must be placed on the Xe atom. This is acceptable because Xe atoms have empty valence shell d orbitals and can accommodate more than eight electrons. The Lewis structure of XeF2 shows two bonding pairs and three lone pairs of electrons around the Xe atom:
        A Lewis diagram shows a xenon atom with three lone pairs of electrons single bonded to two fluorine atoms, each with three lone pairs of electrons.
      2. XeF6: We place three lone pairs of electrons around each F atom, accounting for 36 electrons. Two electrons remain, and this lone pair is placed on the Xe atom:This structure shows a xenon atom single bonded to six fluorine atoms. Each fluorine atom has three lone pairs of electrons.

    Exercise 11.3b

    The halogens form a class of compounds called the interhalogens, in which halogen atoms covalently bond to each other. Write the Lewis structures for the interhalogens BrCl3 and ICl4−.

    Check Your Answer[2]

    Links to Interactive Learning Tools

    Explore Lewis Electron Dot Structures from the Physics Classroom.

    Explore Drawing Lewis Structures from Alchemie.

    Attribution & References

    Except where otherwise noted, this page is adapted JR van Haarlem from “4.4 Lewis Symbols and Structures” In General Chemistry 1 & 2 by Rice University, a derivative of Chemistry (Open Stax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson and is licensed under CC BY 4.0. ​Access for free at Chemistry (OpenStax)​

    1. A Lewis diagram comparing carbon monoxide and carbon dioxide. Carbon monoxide is depicted with both carbon and oxygen each with 1 lone electron pair bonded together via 3 bonds. Carbon dioxide depicts a central carbon atom with two adjacent oxygen atoms. Carbon and oxygen are bonded together via two single bonds between carbon and each oxygen atom. Additionally, oxygen has two lone electron pairs. ↵
    2. Two Lewis structures are shown. The left depicts a bromine atom with two lone pairs of electrons single bonded to three chlorine atoms, each with three lone pairs of electrons. The right shows an iodine atom, with two lone pairs of electrons, single boned to four chlorine atoms, each with three lone pairs of electrons. This structure is surrounded by brackets and has a superscripted negative sign. ↵
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    Enhanced Introductory College Chemistry Copyright © 2023 by Gregory Anderson; Caryn Fahey; Jackie MacDonald; Adrienne Richards; Samantha Sullivan Sauer; J.R. van Haarlem; and David Wegman is licensed under a Creative Commons Attribution 4.0 International License, except where otherwise noted.

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