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DRAWING LEWIS STRUCTURE OF MOLECULES AND POLYATOMIC ATOM

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:

A Lewis structure of calcium is shown. A lone pair of electrons are shown to the right of the symbol.

 Lewis symbols for the elements of the third period 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.
 Lewis symbols illustrating the number of valence electrons for each element in the third period of the periodic table.

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.”

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.”

 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.
 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.

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.

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.

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 Octet Rule. For example , fluorine has an electron configuration notation of. 1s2, 2s2, 2p5. In fluorine’s highest occupied energy level it has 2 electrons in the s sublevel and 5 electrons in the p sublevel. The total number of electrons in fluorine’s highest occupied energy level is therefore = 7. One fluorine atoms bonds to a second fluorine atom to form F2 (fluorine gas, a compound). Each atom shares one of its valence electrons with its partner to get 8 electrons in the valence shell.

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.

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.

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.

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.

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:

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.

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 center.) 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:
      \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}
    • 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):
      \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}
    • 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:
      \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}
    • Since OF2 is a neutral molecule, we simply add the number of valence electrons:
      \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}
  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.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.
  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.
  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.
    • 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.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.
    • In OF2, each atom has an octet as drawn, so nothing changes.

IMPORTANT : no Lewis diagram is complete without formal charges.  Lewis diagrams are drawn to examine mechanisms so knowing which parts of a molecule are electron defficient (+) and which are electron rich (-) is vital.
It is best to have a formal charge of 0 for as many of the atoms in a structure as possible.
If a formal charge of 1- is located next to a formal charge of 1+, the formal charges can usually be minimized by having a lone pair of electrons, located on the atom with the 1- charge become a bonding pair of electrons that is shared with the atom that has the 1+ formal charge (this can be visualised in the same way as the formation of multiple bonds were above).

CAUTION : octets can be expanded to minimise formal charges but only for atoms in the second row of the periodic table  (where n=3 or greater).   For instance in our example, N cannot expand its octet so keeps a formal charge of 1+ and both singly bonded oxygens a formal charge of 1-.  If our molecule were SO3 , however, it would be possible to minimize all formal charges by having the sulfur expand its octet.

LEARNING ABOUT THE DRAWING OF LEWIS STRUCTURE 

 

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