Lesson 38 Lewis Dot Diagrams and Structures

Objectives:

-         The student will draw Lewis Dot Diagrams for various elements.

-         The student will construct Lewis Structures for simple compounds.

-         The student will use Lewis Structures to explain double and triple bonds, and

resonance structures.

 

PA Science and Technology Standards: 3.1.10.B

I. Lewis Structures

a.      Electronegativity differences don’t tell us everything we need to know.

b.      We can use dots to represent valence electrons involved in bonding.

c.      Valence electrons – electrons present in the outermost energy level of an

atom.

d.      Simplest example – hydrogen

                               i.      One hydrogen – H.

                                   ii.      Two together – H:H

Atoms are represented with the number of dots around them that is the same as their group number, if they are in an A group.

f. Examples of single elements:

g.      If they are in a B group (transition metals), they will have a number of electrons determined by the charge they

carry.

h.      Follow the following rules to determine how to make a Lewis structure:

              i.  Determine the total number of valence electrons in the compound

              ii.      Arrange the atoms’ symbols to show how they are bonded and show valence electrons as dots

               iii.      Compare the number of valence electrons used in the structure to the number available from step one

iv. Change to a single dash each pair of dots that represents two shared electrons.

v.    Be sure that all atoms, with the exception of hydrogen, follow the octet rule.

     vi.   General rules to follow:

1.      Hydrogen and halogen atoms usually bond to only one other atom in a

molecule, and are usually on the outside or end of a molecule.

2.      The atom with the smallest electronegativity is often the

central atom.

3.      When a molecule contains more atoms of one element than the others, these atoms often surround the central atom.

vii. Two atoms can share more than one electron pair – try doubling up pairs of electrons if you need more electrons than the original

calculations allow to satisfy the octet rule.

                              

viii. Examples:

i.     A Single bond is the sharing of one pair of electrons between two atoms.

j.      A Lewis structure is a diagram showing the arrangement of valence electrons among the atoms in a molecule.

II. Two atoms can share more than one electron pair

a. Carbon is a good example of an atom that can share more than one pair of electrons -it

can share two or three with another carbon atom

b.      Several other elements, such as nitrogen, oxygen, and sulfur, can also share more than one pair.

                                          i.      A double bond is a covalent bond formed by the sharing of two

pairs of electrons between two atoms                                          ii.      A triple bond is a covalent bond

formed by the sharing of three pairs of electrons between two atoms.

c.      If no arrangement of single bonds provides an appropriate Lewis structure, consider looking into forming multiple bonds.

d.      Examples:

a.      Examples:

Lewis Structures Virtual Lab 2 Video

III. Sometimes no single Lewis structure is correct

a.      For some molecules, it is not possible to draw one correct Lewis structure.

b.      NO2 is a good example of this. Example:

 

c. Things to think about

             i. Atoms with unpaired electrons are somewhat unstable compared to atoms with all

electrons paired, but they are possible.

               ii.      Situations where an octet does not form are normally unstable, but possible in some

cases also.

               iii.      Before such a rule breaking structure can be considered accurate, either an octet

must be impossible, or a non-octet structure must be found to be experimentally more stable.

               iv.      Example

d.      Often several equivalent structures can be drawn – These are called resonance hybrids.

             i.      Originally, it was thought that the structures resonated between the possible

structures

             ii.    It was later determined that the actual structures are an average, or combination, of these structures.

             iii.      Examples:

IV. Lewis structures can be drawn for polyatomic ions.

a.      Polyatomic ions are held together by covalent bonds – we can draw Lewis structures for them just as

easily.

b.      Ammonium – NH4

                       i.  The ion is prepared from ammonia, NH3, and a

proton (the nucleus of a hydrogen atom)

                       ii.  The hydrogen shares a pair of electrons donated by the nitrogen

                      iii. This produces a coordinate covalent bond – a bond where two atoms share a pair of

electrons donated by only one of the atoms.

 iv. Once it is formed, this coordinate covalent bond is indistinguishable from a normal covalent bond.

  v. Brackets are placed around the structure, with the charge on the ion shown outside the brackets.

vi. When you count up electrons that should be in the structure, you need to account for the

charge on the ion.

1. If it is a positive charge, subtract that number of electrons from the

total.

2.    If it is a negative charge, add that number of electrons to the total.

   vii.      Examples:

Questions: 1. Create Lewis dot diagrams for each of the following

elements:

a. Nab. Hec. Sd. Pe. Mgf. Ig. Oh. Ali. Ne

1. Create Lewis dot diagrams for each of the following compounds:

a. HI

b. O2

c. CCl4d. SO2

e. ICl

f. Cl2g. H2O

h. CH3COOH

i. NO3-

j. SO4-2

k. PO4-3

l. N2

CO32- HCPNBr3 O2

SiO2 P2

BI3 HCBrF2

ArO2 CS2

AlBr3 GeI4