Dot Structures of Molecules Laboratory Exercise:


After completing the Data Page for the Atomic Spectra Lab, adjourn to the classroom to work on Drawing Lewis Dot Structures of Molecules Lab.

You will need scissors to cut out the paper atom pages provided by the instructor. Keep the paper atoms in an envelop. A seven page Data packet has been provided for you to directly draw the 33 compounds in both dot and stick structure for your lab report. This report is due in one week.

In one week, the instructor will provided you with an answer packet for you to self check your work during the next lab before you submit it. This exercise is worth up to 30 points depending on your instructor.

CHM 1025C Text Reference:

The current CHM 1025C textbook: An Introduction to Chemistry (Atoms First) , the author, Mark Bishop has Table 5.1 page 195 demonstrating the stick structures of various nonmetals. Read sections 5.4 and 5.5 (Drawing Lewis Structures), especially the General Procedures on pages 198-199-200. Bishop describes seven steps on pages 198-199 for drawing Lewis Dot Structures for molecules.

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 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:


If you still have the old Corwin CHM 1025C text, (Introductory Chemistry-Concepts & Connections – 5th Edition) sections 12.4 and 12.5 beginning on page 330 has a simpler discussion. On the bottom of page 334 Corwin has only four rules similar to above.

CHM 2045C Text References:

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 this lab 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 encountered 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.

     CHM 1020 Text References:

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:

John Taylor’s Method for Drawing Dot Structures

On John Taylor’s web site he has a lengthy study guide for Polyatomic ions:  

From that study guide he modifies the mathematical approach given in each of the books on the previous pages. The following are his seven steps:

Using your envelope of paper atoms, assemble a reasonable dot structure of each polyatomic ion in the molecular form of acids (not in water) using the following criteria (You may print out the following pages and cut out the atoms: O, H, S and C, Cl, P, N ):

  1. Place the nonmetal which is not oxygen or hydrogen in the middle of your desk.
  2. First using all single bonds, hook all oxygens in the formula to the central nonmetal (Simple covalent or coordinate covalent bonds).

     Never hook oxygen to oxygen except in peroxides, O2, O3.

  1. If oxygen is present, hook the hydrogen written first in the chemical formula to an oxygen. Hydrogen requires only two electrons to fill it's orbital. Notice the change in the polyion's charge when a hydrogen is placed on the oxygen. If hydrogen is written second in the formula after another nonmetal, then hook that hydrogen to that nonmetal, not to oxygen such as (CH3-COOH written organically) acetic acid HC2H3O2 or oxalic acid H2C2O4 both hydrogens are attached to the oxygens.
  2. 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.
  2. Extras: 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. Never have more than two bonds to any oxygen (except CO). If you place hydrogen to an oxygen, then oxygen HAS to hook to another element by a single bond, never a double bond.


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: