Module 4 Part B: Dot Structures of Molecules Study Notes:

For CHM 1025C Students:

In the 6th Edition of Corwin, Section 12.1 discusses the Chemical Bond Concept. The Octet Rule is explained. The two ways the octet rule can be achieved is by either electron transfer or by sharing one or more pairs of electrons (Hydrogen follows the Duet rule or rule of two).

Section 12.2 discusses the Ionic Bond, the transfer of electrons and the formula unit concept.

A fundamental particle held together by covalent bonds is a molecule.  Section 12.3 discusses the covalent Bond. Also there is a discussion of the Bond Length and Bond Energy.

In Section 12.4 is the discussion of the Electron Dot Formulas of Molecules. In Section 6.8 we drew electron dot formulas of atoms and pretested them in M-3 Section D:

         The number of dots around each atom is equal to the number of valence electrons the atom has. (The Roman Numeral A groups give us that number, except Helium)

         We will now draw electron dot formulas for molecules (also called Lewis structures).

         A Lewis structure shows the bonds between atoms and helps us visualize the arrangement of atoms in a molecule. (A two dimensional model)

Guidelines for Electron Dot Formulas of Molecules (Section 12.4)

1.      Calculate the total number of valence electrons by adding all of the valence electrons for each atom in the molecule.

2.      Divide the total valence electrons by 2 to find the number of electron pairs in the molecule.

3.      Surround the central atom with four electron pairs. Use the remaining electron pairs to complete the octet around the other atoms. The only exception is hydrogen, which only needs two electrons.

4.      Electron pairs that are shared by atoms are called bonding electrons. The other electrons complete octets and are called nonbonding electrons, or lone pairs or unshared pairs.

5.      If there are not enough electron pairs to provide each atom with an octet, move a nonbonding electron pair between two atoms that already share an electron pair. (Creates a double covalent bond.)

Electron Dot Formula for H2O (section 12.4)

1.      Count the total number of valence electrons: oxygen has six and each hydrogen has one for a total of eight electrons [6 + 2(1) = 8 e-]. The number of electron pairs is 4 (8/2 = 4).
 

2.      Place eight electrons around the central oxygen atom.

 12_Pg331_UnFigure_2

3.      We can then place the two hydrogen atoms in any of the four electron pair positions.  Notice there are two bonding and two nonbonding electron pairs.

12_Pg331_UnFigure_3

4.      To simplify, we represent bonding electron pairs with a single dash line called a single bond.

5.      The resulting structure is referred to as the structural formula of the molecule (Dot Stick model of the molecule).

6.      Remember it doesn’t matter which two electron pairs the two hydrogen atoms are placed as this is a two dimensional picture of what is a three dimensional molecule. Next week’s lab will focus on the three dimensional Structural Formulas.


7.       The double bond is explained in Section 12.4 by drawing the Lewis Structure for SO3 (Sulfur Trioxide)

     or   

8.      The Triple Bond is explained in Section 12.4 by drawing the Lewis Dot Structure of Hydrogen Cyanide (HCN):

           or       

9.      The Structural Formula or Dot Stick Model is explained in Section 12.4 via H2O; SO3; and HCN as shown above.

10.  Section 12.8 explains the concept of the coordinate covalent bond (One atom provided both shared pair of electrons to make the covalent bond.) The example is ozone (O3) built from Oxygen gas (O2):

12_Pg342_UnFigure_2

John Taylor’s Method for Drawing Dot Structures

On our course web site I have several interactive and static pages for discovery of the formula, charge, and naming of Polyatomic ions:

Module 4: Polyatomic Ions:

Intro To Polyatomic Ions
New Intro 2/10/09

Interactive Periodic Chart:
Polyatomic Ions
New Web Site: Updated 6/9/11

Memorizing Polyatomic Ions:
The Problem


Discover Formulas and Charges:
NonMetal Polyions (Box of 4)

Discover Formulas and Charges:
NonMetal Polyions (Box of 6)

Discover Formulas and Charges:
Transitional Metal Polyions

The NonMetal -ate Answers:
6ChargeBoth 6FormulaBoth 6Both

Laurenzo's Nonmetal -ate Screen

The Transitional Metal -ate Answers:
TChargeAll TFormulaBoth TBoth

Laurenzo's Transitional Metal -ate Screen

193rd 2YC3 Conference (September 17, 2011) Abstract:
How to Teach Polyatomic Ions in Chemistry

Taylor's -ate 3/4 Oxygen Rule

Taylor's -ate Charge Rule

Discover Formulas and Charges:
-ite Rule Polyatomic Ions

Formulas and Charges: -ite Ions

Formulas and Charges: -ide ions

The Complete Polyatomice Ions List

Dot Structures using:
Simple Octet Rule!

Nonmetal polyatomic Ions 

 

 

 

From these web pages I have the following steps which I use the first day of class when we begin drawing dot structures of molecules:

Create an envelope of paper atoms, assemble a reasonable dot structure of each polyatomic ion of molecular acids (not in water) using the following criteria

  1. Dot/Stick Structures of Atoms:
    Oxygen-Carbon Dot/Stick Atoms
    Hydrogen-Chlorine-Nitrogen Dot/Stick Atoms
    Hydrogen-Phosphorus-Sulfur Dot/Stick Atoms
    Oxygen-Hydrogen-Carbon-Chlorine Dot Stick Atoms

    Dot Structures of Atoms:
    O, H, S, e-1 atoms
    P, N, Cl, e-1 atoms
  2. Place the nonmetal which is not oxygen or hydrogen in the middle of the structure.
  3. If oxygen and hydrogen are both present, hook the hydrogen written first in the chemical formula to an oxygen atom. Hydrogen requires only two electrons to fill it's orbital (duet rule or rule of two)l. If hydrogen is written second in the formula after another nonmetal, then hook those hydrogen atoms to that nonmetal, not to oxygen such as (CH3-COOH written organically) acetic acid HC2H3O2 while in oxalic acid H2C2O4 both hydrogen atoms are attached to the oxygen atoms.
  4. If using all single bonds connecting the oxygen to the central nonmetal, the count of electrons around each element should total eight electrons (octet rule). This includes the element's original outer surface (valence) electrons plus the electrons being shared from the bonded element.

 
                    


 

 

  1. If the count is 7-7, then add a second bond, a double covalent bond (four electrons being shared between two atoms.

                                  

                      

  1. If the count is 6-6, then make a triple covalent bond between the two elements (six electrons shared between the two atoms).
                  

 

  1. If the count is 8-6, (or 9-7) then make a coordinate covalent bond. Hook the six (vacant) orbital onto an unshared pair of the eight. A coordinate covalent bond is still a single bond. In making double, triple bonds you may also use one or two coordinate covalent bonds to predict a structure using the octet rule, if necessary as in carbon monoxide.

 

 

 

 

H.   Extras:

1.  Never have an unshared pair or lone pair of electrons on a carbon atom (except carbon monoxide, CO). These are two dimensional structures, so there are many variations of the answer shown on the web site.

2.   Never have more than two bonds to any oxygen (except CO). If you place a hydrogen atom to an oxygen atom, then oxygen has to hook to another element by a single bond, never a double bond.

3.  Nitrogen should have at least three bonds in a molecular structure: three single bonds, or a double and single bond, or a triple bond.

4.  Carbon should always have four bonds: four single bonds; or one double and two single bonds; or two double bonds; or one triple bone and one single bind. See the chart below:

 

From the Kotz Textbook:

In Section 9.4 of the 6th edition of Kotz (CHM 2045C) there is an excellent explanation of covalent bonds and drawing dot structures of molecules. We will focus in Part B only on those molecules which can be explained by the octet rule for the nonmetals (and duet rule for hydrogen). Page 384-5 lists five steps where you do a little math in step 2 to calculate all the valence electrons and the number of bonds and lone pairs. Examples 9.2 and 9.3 show simple binary molecular compounds. Please note the Problem-Solving tip on page 388.  Tables 9.4 and 9.5 show the comparison of molecules and polyatomic ions.

Kotz’s five steps (#2 was expanded to Steps 2 and 3 below):

1.    Decide on the central atom (usually not oxygen or hydrogen). The central atom is usually the one with the lowest electron affinity. In formaldehyde, CH2O, the central atom is carbon.

2.    Determine the total number of valence electrons in the molecules or the ion. I a neutral molecule this number will be the sum of the valence electrons of each atom. For a negative ion add the charge to this total number. For a positive ion subtract the positive charge from this total number. For CH2O:  C=4, H=1(x2) , O=6 this would be 4+2+6 = 12 total valence electrons.

3.    Take this total number of electrons and divide by two to determine the number of electron pairs. For CH2O:  12/2 = 6 electron pairs

4.    Place one pair of electrons between each pair of bonded atoms to form a single bond. You can either show a pair of dots, or draw a single stick between the two atoms to represent the single covalent bond.

5.    Use any remaining pairs as lone pairs around each atom (except hydrogen) so that each atom is surrounded by eight electrons. (There is never a lone pair on a carbon except in Carbon moxide)

6.    If the central atom has fewer than eight electrons at this point, move one or more of the lone pairs on the terminal atoms in a position intermediate between the center and the terminal atom to form  multiple bonds. (As a general rule double or triple bonds are formed when both atoms are from the following nonmetals: C, N, O, S. That is, bonds such as C=C, C=n, C=O, S=O will be incountered frequently.

In the Kotz 7th edition Chapter 9 has become Chapter 8 as the two atomic theory chapters 7& 8 have been combined into one chapter 7. The five steps are on pages 353-4 and example 8.1 should be studies on pages 354-355.

Other textbooks:

McMurray’s Text (CHM 2045C) has Covalent Bonds and Molecular structure in Chapter 7. In section 7.5 McMurray summaries the dot structures of compounds of the nonmetals in the second row of the periodic table pages 229-232. The mathematical process similar to Kotz’s six steps above is found on pages 235-236.

Brady’s text (CHM 2045C) covers covalent bonding in Chapter 9: Chemical Bonding-General Concepts. Brady covers drawing dot/stick structures of molecules in section 9.7 On page 377, the six step similar to the above are listed.

In the CHM 1020 text, Hill, (Chemistry for Changing Times-11th edition) Chapter 5 discusses chemical bonds.  Covalent bonds are introduced in section 5.7. However, section 5.11(Rules for Writing Lewis Formulas) on pages 132 and 133 list the five general rules for drawing Lewis Dot Structures. Table 5.5 on page 135 is a good summary which is sown on a previous page:

CHM 1032C Text Reference:

The current CHM 1032C textbook: Fundamentals of General, General Organic, and Biological Chemistry; 6th edition; the authors cover Molecular compounds in Chapter 5. Section 5.5 Molecular Formulas and Lewis Structures provide you with the general directions for drawing Lewis Dot Structures. On page 119 they break down the process in five steps.

Table 9.4 from Kotz displays Common Hydrogen containing Compounds and Ions:

 

 

 

 

 

 

Table 9.5 from Kotz displays common oxoacids and their anions:

 

 

 

 

 

From McMurray Stick Structures:

 

 

 

 

 

Table 5.1 from Bishop Text:

Summary of Bishop’s Seven Steps to Draw Lewis Dot Structures:

Bishop’s Seven Steps to Draw Lewis Dot Structures:



                  
                  

 

 

CHM 2045C Text Reference Bonding Flow Chart: