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)
the total number of valence electrons by adding all of the valence electrons
for each atom in the molecule.
total valence electrons by 2 to find the number of electron pairs in 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)
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.
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.
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).
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)
8. The Triple Bond is explained in Section 12.4 by drawing the Lewis Dot Structure of Hydrogen Cyanide (HCN):
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):
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:
New Web Site: Updated 6/9/11
Memorizing Polyatomic Ions:
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. 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.
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):
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.
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.
Take this total number of
electrons and divide by two to determine the number of electron pairs. For CH2O: 12/2 = 6 electron
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.
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
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: