Bonding :  Lewis Dot Structures

Lewis Dot Structure - A symbolic description of the distribution of valence electrons in a molecule.  Dots are used to represent individual electrons and lines are used to represent covalent bonds.

 Drawing Lewis Dot Structures Add up the total number of valence electrons in the molecule by totalling the valence electrons on each atom in the molecule or polyatomic ion. If you are drawing a Lewis Structure for a polyatomic ion: Negative Ion:  Add the number of electrons equal to the negative charge on the ion. Positive Ion:  Subtract the number of electrons equal to the positive charge on the ion. Draw the skeleton structure of the molecule or polyatomic ion in which the covalent bonds between the atoms are drawn as single lines.  Each bond equals two valence electrons.  If the molecule has more than two atoms, the atom with the lowest electronegativity is generally the central atom and is written in the middle. Distribute valence electrons around the outer atoms as nonbonding electrons to until each atom has a complete outer shell (i.e. 8 electrons except for H which has only 2 valence electrons). Add the remaining valence electrons to the central atom. BONDING ELECTRONS + NONBONDING ELECTRONS = TOTAL VALENCE ELECTRONS Check the central atom. If the central atom has eight electrons surrounding it, the Lewis Structure is complete. If the central atom has less than eight electrons, remove a nonbonding electron pair from one of the outer atoms and for a double bond between that atom and the central atom.  If needed, continue to remove nonbonding electron pairs from the outer atoms until the central atom has a complete octet. If the central atom has more than eight electrons, then this means that the central atom has expanded its valence shell to hold more than eight electrons.  This is allowed for atoms with valence shells in the third energy level or higher (i.e. in or beyond the third period of the periodic table).

How can a central atom have more than 8 electrons in its valence shell?

 [Ne] 3s23p33d0
• Since the valence shell is in the third energy level, there is a whole set of d orbitals that can hold more electrons.  This P atom, therefore, can expand its outer shell to hold more valence electrons.
• Remember that atoms in the second energy level or lower cannot expand their valence shells because 1d and 2d orbitals do not exist.
• Si can expand its valence shell whereas C cannot.
• P can expand its valence shell whereas N cannot.
• S can expand its valence shell whereas O cannot.
• Cl can expand its valence shell whereas F cannot.

Example Lewis Dot Structures

1) Add up the total number of electrons:

2) Draw the skeleton structure:

3) Distribute the valence electrons around the atoms until each atom has a complete valence shell (eight valence electrons).

This Lewis Structure is complete.  If we check to make sure all the electrons have been accounted for, we find there are 14 electrons (with the single covalent bond counting as two bonding electrons).

1) Add up the total number of electrons:

2) Draw the skeleton structure:

4) Since H atoms only have 2 valence electrons, each H in this skeleton structure has a complete outer shell.  Since all 8 electrons have been used, we double check to make sure that the central atom (C) has eight surrounding electrons, which it does.  This Lewis Structure is complete.

1) Add up the total number of electrons:

2) Draw the skeleton structure:

3) Since H atoms only have 2 valence electrons, each H already has a complete outer shell.  We add the remaining 2 valence electrons to the central atom (N).

4) Check the central atom to make sure it has a complete octet.  It does, so the Lewis Structure is complete.

1) Add up the total number of electrons:

2) Draw the skeleton structure:

3) Since H atoms only have 2 valence electrons, each H already has a complete outer shell.  All of the valence electrons have been used, so we double check to make sure that the central atom (N) is surrounded by eight electrons.  It is, so the Lewis Structure is complete.

There is, however, one thing missing.  The molecule is a polyatomic ion, but the Lewis Structure thusfar does not reflect the +1 charge on the polyatomic ion.  This introduces the concept of formal charge.
 Formal Charge - The charge on an individual atom in a Lewis Structure.  To calculate the formal charge, we must do the following. Divide each covalent bond down the middle, distributing one electron from the two bonding electrons in the covalent bond to each atom that formed that bond. Add up the electrons directly surrounding each atom after this division has been done. Compare the total number of electrons around each individual atom in the Lewis Structure to the number of valence electrons in each respective neutral atom. If the atom in the Lewis Structure has more electrons than its respective neutral element, the formal charge on that atom is negative.  The number of the formal charge is equal to the difference between the electrons around the atom in the Lewis Structure and the valence electrons of the respective neutral atom. If the atom in the Lewis Structure has fewer electrons than its respective neutral element, the formal charge on that atom is positive.  The number of the formal charge is equal to the difference between the valence electrons of the respective neutral atom and the electrons around the atom in the Lewis Structure. Find the net formal charge (i.e. formal charge of the overall molecule or ion) by adding up all the formal charges of each atom in the Lewis Structure.  The net formal charge should be the same as the charge on the molecule or ion.  Therefore, neutral molecules should have a net formal charge of 0.

So let's return to our last Lewis Structure and calculate the formal charge on each of the atoms.

1) Divide the covalent bonds and distribute the electrons:

2) Add up the electrons surrounding each atom.

• Each H atom has 1 electron.  The central atom (N) has 4 electrons.

3) Compare the total number of electrons around each individual atom in the Lewis Structure to the number of valence electrons in each respective neutral atom.

• A neutral H atom has 1 valence electron and so do the H atoms in the Lewis Structure.  Therefore the formal charge on all of the H atoms is 0.
• A neutral N atom has 5 valence electrons, which is one more than the N atom in the Lewis Structure.  Therefore, the formal charge on the N atom is +1.

4) Find the net formal charge by adding up all the formal charges of each atom in the Lewis Structure.

• The only atom with a formal charge was the N atom, so the overall formal charge of the Lewis Structure is +1 which is correct because the charge on NH4+ is +1.

This is the final Lewis Structure for NH4+.

1) Add up the total number of electrons:

2) Draw the skeleton structure:

3) Distribute the valence electrons around the atoms until each atom has a complete valence shell (eight valence electrons).

4) Since all 32 electrons have been used, we double check the central atom (P) to make sure that it has a complete octet.  It does, so now we calculate the formal charge on the Lewis Structure.  (See above notes on formal charge)

1) Divide the covalent bonds and distribute the electrons:

2) Add up the electrons surrounding each atom.

• Each O atom has 7 electrons.  The central atom (P) has 4 electrons.

3) Compare the total number of electrons around each individual atom in the Lewis Structure to the number of valence electrons in each respective neutral atom.

• A neutral O atom has 6 valence electrons, which is one less than the O atoms in the Lewis Structure.  Therefore, the formal charge on each of the O atoms is -1.
• A neutral P atom has 5 valence electrons, which is one more than the N atom in the Lewis Structure.  Therefore, the formal charge on the N atom is +1.

4) Find the net formal charge by adding up all the formal charges of each atom in the Lewis Structure.

• The four O atoms each have a -1 FC totalling -4.  The P atom has a +1 FC.  The net FC is -3 which is correct because the charge on PO43- is -3.

1) Add up the total number of electrons:

2) Draw the skeleton structure:

3) Distribute the valence electrons around the atoms until each atom has a complete valence shell (eight valence electrons):

4) Since all 16 electrons have been used, we double check the central atom (C) to make sure that it has a complete octet.  The C only has 4 surrounding electrons with the two single bonds on both sides.  Therefore, we have to remove one pair of nonbonding electrons from one of the O atoms and form a double bond between the C and that O atom.

5) We check again, and the C atom now has 6 surrounding electrons.  This is still too few, so we have to remove another pair of nonbonding electrons from the other O atom and form another double bond between the C and the other O atom.

6) Now the central atom (C) has eight surrounding electrons.  Both of the O atoms also have eight surrounding electrons, so the Lewis Structure is complete.  If we check the formal charge on each of the atoms, C has a 0 FC and so do both O atoms, so the net FC is zero and does not have to be noted.

1) Add up the total number of electrons:

2) Draw the skeleton structure:

3) Distribute the valence electrons around the atoms until each atom has a complete valence shell (eight valence electrons):

4) Add the remaining electrons to the central atom (S):

5) Check to make sure that the central atom has a complete octet.  In this case, the S atom has only 6 of the 8 required electrons.  Therefore, we must remove one of the nonbonding electrons around one of the O atoms and create a double bond between that O atom and the S atom:

This, however, raises one question.  From which O atom do we remove the nonbonding electrons to form the double bond.  The two structures below are basically the same except for the choice of the oxygen atom with which the S atom forms the double bond.

Actually, both of these structures are incorrect.  They both suggest that one bond (the double bond) in the molecule is stronger than another (the single bond) which is untrue.  The two S-O bonds are equally strong.  This introduces the concept of resonance hybrids.

• Resonance hybrid - A mixture or average of two or more Lewis Structures.

To indicate that SO2 is a resonance hybrid of the two Lewis Structures above, a forward/backward arrow is drawn between the two Lewis Structures.

When the formal charge of each atom in this Lewis Structure is checked, the one O atom with 3 pairs of nonbonding electrons has a -1 FC, and the S atom has a +1 FC.  The net FC is therefore zero and does not need to be indicated.

Next:  "Predicting Molecular Shapes - VSEPR"