Lewis Diagrams Molecular Geometry Worksheet

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In covalent compounds, electrons are shared between atoms. Lewis structures are often used to represent covalent bonding. Atoms in covalent compounds will form enough bonds to obtain eight electrons in their valence shell, a property known as the octet rule. There are exceptions to the octet rule – one such exception is called the duet rule which occurs with hydrogen (H) atoms. Because only two electrons fit into the n = 1 shell, H atoms can have a maximum of a single bond, or 2 electrons total. There are also atoms in compounds that form incomplete octets (less than 8 electrons) or expanded octets (more than 8 electrons).


Lewis structures typically contain shared pairs of electrons, or bonds, and lone pairs of electrons. (PIC?) The number of bonds that form between atoms in a compound is directly related to the number of valence electrons in a compound. The number of valence electrons for a compound is always equal to the total number of valence electrons for its atoms. For a main group element, the number of valence electrons for an atom is equal to its group number. For example, oxygen (O), is in Group VIA and has 6 valence electrons, and Carbon (C), is in Group IVA and has 4 valence electrons. Therefore carbon monoxide (CO) has a total of 6+4, or 10 valence electrons.


Because carbon monoxide does not contain enough valence electrons to fill both the carbon’s and the oxygen’s valence shells separately, some of the 10 electrons will be share by both the carbon and the oxygen. A single bond occurs between atoms, when 2 electrons are shared between two atoms, meaning the valence shells of the two atoms overlap so the electrons can exist in the valence shell of both atoms. A double bond occurs between two atoms when 4 electrons are shared. A triple bond occurs between two atoms when 6 electrons are shared in a bond. Electrons that are not shared between two atoms are called lone pair electrons. These electrons contribute to the total valence electron count for the compound. The Lewis structures of the common compounds, ammonia, water, and hydrogen fluoride are shown in Figure 14.1 below. These structures contain only single bonds. The Lewis structures for compounds with double and triple bonds with lone pairs are shown in Figures 14.2 and 14.3 below.


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Figure 14.1 – Lewis Structures of Common Compounds containing Single Bonds



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Figure 14.2 – Lewis Structures of Common Compounds containing Double and Triple Bonds

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To draw the Lewis structure for a compound, follow the steps below:


  1. Determine the number of valence electrons for the compound. For cations, subtract an electron for each positive charge and for anions, add one electron for each negative charge.
  2. Draw a “skeleton” structure for each molecule or ion, arranging the atoms around the central atom, which is generally the least electronegative atom in the compound.
  3. Connect each atom to the central atom with a single bond (one electron pair).
  4. Distribute the remaining electrons as lone pairs on the terminal atoms, completing an octet around each atom. (Remember that H atoms only have two electrons to fill the valence shell).
  5. Place all remaining electrons on the central atom.
  6. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom to obtain octets when needed.


Lewis structures simply show the linkages between atoms and the presence of lone pairs. They do not, by themselves, show the three-dimensional arrangement of atoms in space. The Valence Shell Electron Pair Repulsion (VSEPR) theory develops Lewis’s ideas so that we can predict the shapes of simple molecules. The VSEPR theory adds rules that account for bond angles.


Rule 1: Regions of high electron concentration (bonds and lone pairs on the central atom) repel one another and, to minimize the repulsions, these regions move as far as possible from each other while maintaining the same distance from the central atom.

Rule 2: There is no distinction between single and multiple bonds: a multiple bond is treated as a single region of high electron concentration.

Rule 3: All regions of high electron concentration, lone pairs and bonds, are included in a description of the electronic arrangement, but only the positions of atoms are considered when identifying the shape of a molecule.


We begin by looking at molecules that consist of one central atom to which all the other atoms are attached with no lone pairs on the central atom. (Fig. 1) The molecular shape is the same as the electron arrangement in these molecules.


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Figure 14.3 The names of the shapes of simple molecules with no lone pairs on the central atom.



The bond angles, the angles between the bonds, are fixed by the symmetry of the molecules as shown in Figure 14.3: linear (180O), trigonal planar (120O), and tetrahedral (109.5O). When there is more than one central atom in a molecule, the molecular geometry can be determined on each central atom.


Now we consider molecules with one or more lone pairs on the central atom. (Figure 14.4) If lone pairs are present, the molecular shape differs from the electron arrangement because only the positions of the atoms are considered when naming the shape.



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Figure 14.4 – The names of the shapes of some simple molecules with lone pairs (located on top of the central atom) on the central atom.


The electrons in the molecules shown here are arranged in a tetrahedral geometry, but have a different molecular shape. The presence of lone pairs on the central atom makes distinction between the electron geometry and the molecular shape. To help predict the shapes of molecules, we use the generic “VSEPR formula”: AXnEm, where A represents a central atom, X represents an attached atom, and E represents a lone pair on the central atom.

Table 14.1 Molecular Shapes Predicted by VSEPR



(Electronic geometry)


cular geometry with VSEPR for


0 lone pair

1 lone pair

2 lone pairs






Linear (AX2)





Trigonal planar




Trigonal planar (AX3)

< 120O


Bent (AX2E)









< 109O

Trigonal Pyramid (AX3E)

< 109O Bent




In covalent bonds, electrons are shared between atoms. However, the electrons in a covalent bond are not always shared evenly. A polar covalent bond results from the uneven sharing of electrons between two atoms, a polar covalent bond results. A polar covalent bond is characterized by a partial positive charge (δ+) and a partial negative charge (δ) on opposite ends of the bond. A polar molecule is a molecule that displays a partial positive and partial negative charge on opposite ends of the molecule. Polarity of a molecule arises from two factors: (1) the presence of a polar covalent bond within the molecule and (2) the shape of the molecule. We can determine the polarity of a molecule in two different cases:


Case 1: A diatomic molecule*

  1. A diatomic molecule is polar if its bond is polar. (Ex) An HCl molecule: a polar molecule with its polar covalent bond (δ+H−Clδ−).


All diatomic molecules composed of atoms of two different elements are at least slightly polar.


  1. A homonuclear diatomic molecule, a diatomic molecule built from two atoms of the same element, such as O2, N2, and Cl2, is nonpolar, because its bond is nonpolar.


Case 2: A polyatomic molecule*


  1. A polyatomic molecule may be nonpolar even if its bonds are polar. 180O



    1. CO2 is nonpolar: the two δ+C=Oδ− dipole moments in carbon dioxide, a linear molecule, point an opposite direction, and so they cancel each other.


    1. CCl4 is nonpolar: If the four atoms attached to the central atom in a tetrahedral molecule are the same, the polar bonds cancel and the molecule is nonpolar.




  1. A polyatomic molecule may be polar if its bonds are polar and they do not cancel each other.


    1. H2O is polar: the two δ+H−Oδ− dipole moments in H2O lie at 104.5O to each other and do not cancel. This polarity explains why water is such a good solvent for ionic compounds.

< 109O


    1. Both CHCl3 and NH3 are polar: If one or more of the atoms are replaced by different atoms (as in CHCl3) or by lone pairs (as in NH3), then the polarity associated with the bonds are not the same, so they do not cancel.






< 109O

*Lone pairs on terminal atoms are not shown in the VSEPR structures, because they are not included when identifying molecular shapes.





  1. Fill in the Data Table (total valence electrons and Lewis Diagram) for each of the first four compounds in the table.


Step 1: Draw the Lewis structure.

Step 2: Assign the VSEPR electron geometry

Step 3: Assign the VSEPR molecular geometry.

Step 4: Identify the shape considering only atoms. Determine if the polar bonds cancel.


  1. Go to https://phet.colorado.edu/sims/html/molecule-shapes/latest/molecule-shapes_en.html 3. Choose Model.
  2. To the central atom, add the number of lone pairs, singly bonded, doubly bonded, and triply bonded atoms that you predicted in the Lewis structure.
  3. At the bottom left, you can check the boxes to reveal the electron geometry and the molecular geometry that corresponds to the predicted bonding. You can also rotate the molecule by clicking and dragging it. This may help you to visualize each geometry. You can verify your predictions with these molecules.
  4. Models of some of the molecules listed in the Data Table are included in the Real Molecules portion of the simulation. You can verify your predictions with these molecules.
  5. Have your instructor check your predictions in the Data Table.
  6. Complete the Data Table.







Lewis Structure and Electron Geometries






Lewis Diagram

Electron Geometry

Molecular Geometry

Polar or Not?

Graded By


































Lewis Diagram

Electron Geometry

Molecular Geometry

Polar or Not?

Graded By Instructor

































Lewis Diagram

Electron Geometry

Molecular Geometry

Polar or Not?

Graded By Instructor
















(Hint: O in the terminal position)

















  1. Which of the molecules in your Data Sheet have the molecular shape different from the electron group arrangement around the central atom? Why?












  1. CCl4 is a nonpolar molecule, while CHCl3 and CH2Cl2 are polar molecules. Draw the Lewis structures of these three molecules. Explain the observation in polarity of the molecules.












  1. Define resonance. Which of the covalent compounds from today’s experiments show resonance. Draw all of the resonance structures for the compounds.













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  1. The VSEPR model extends Lewis’s theory to account for molecular shapes. Write the rules of the VSEPR model that account for molecular shapes and bond angles:
        1. Rule 1



        1. Rule 2



        1. Rule 3



  1. The VSEPR formula, AXnEm, helps us to predict the molecular shape. What does each symbol represent in the formula?
      1. A:



      1. X:



      1. E:



  1. Select a molecule in which the molecular shape is the same with the electron arrangement. Explain your reasoning.


    1. CO2 (b) H2O (c) NH3 (d) SO2









  1. Select a nonpolar molecule in which the dipole moment of polar covalent bonds cancels each other. Explain your reasoning.


    1. H2O (b) NH3 (c) CCl4 (d) HCl





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