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Chemistry 121: Appendix C
Laboratory Tutorials
Contents of this section:
Covalent Bonding Lewis Structures Resonance Structures Valence Shell Electron Pair Repulsion Theory (VSEPR) Polarity Drawing Molecules in 3-D Drawing abbreviated molecular structures
Covalent Bonding
Lewis structures are simple two-dimensional drawings of three-dimensional molecules. They allow chemists to quickly and easily determine the major characteristics of almost any covalently-bonded molecule, including the presence or absence of multiple bonds and lone pair electrons, the orientation and angle of most atoms in a molecule, the overall molecular shape, and finally the molecular polarity, including the relative size and direction of the electrical dipole (if present).
Guidelines for Drawing Lewis Structures
- Add up the total number of valence electrons for all the atoms in the molecule.
- Write the element symbol for the central atom (usually the least electronegative). Note that H and F must always be outer atoms since each atom only forms one bond.
- Write the element symbols for all the outer atoms around the central atom. Draw a straight line to connect each outer atom to the central atom.
- Distribute the remaining valence electrons in pairs to make an octet for all atoms. Start first with the outer atoms and then finish with the central atom.
- If the central atom still does not have an octet, move a lone pair of electrons from an outer atom to be shared with the central atom to make a multiple bond.
Examples of Lewis Structures
Case 1: Draw the Lewis structure for CH 3 Cl.
A. Calculate the total number of valence electrons for all atoms in the compound:
1 carbon (4e) + 3 hydrogen (1e each, total of 3e) + 1 chlorine (7e) = 14 e–^ or 7 pairs.
B. Choose C to be central, since it is the least electronegative (most willing to share electrons with many outer atoms). Although H is actually less electronegative than C, H can only form one bond, and must always be an outer atom.
C. Connect all the atoms with single bonds (see below). This uses up 8 electrons (4 pairs).
D. Put the last three pairs of electrons around Cl, since the C atom has an octet, and each H atom has a pair of electrons.
Finally, confirm that electrons were neither created nor destroyed – that there are in fact 14 electrons, no more or less, in the final Lewis structure.
The Lewis electron-dot structure for CH 3 Cl is shown below.
Figure 1. Drawing the Lewis structure for CH 3 Cl
Case 2: Draw the Lewis electron-dot structure for nitrite ion, NO 2 –.
A. Calculate the total number of valence electrons for all atoms in the compound:
1 nitrogen (5e) + 2 oxygens (6e each, total of 12e) + 1e from the charge on the ion = 18e or 9 pairs of electrons.
B. The central atom is usually the least electronegative, N is the central atom, and the O atoms are the outer atoms. Connect all the atoms with single bonds. This uses two pairs of electrons.
C. Distribute the remaining seven electron pairs around the outer atoms, then put the last pair on the N atom (see Figure 2 Step 4 below).
Figure 2: Drawing the Lewis structure for NO2-
D. While the two O atoms have an octet, the N atom needs one more pair. Giving a lone pair of electrons from an O atom to the N atom would prevent the O atom from having an octet. For all the atoms to get an octet, one O atom must share a lone pair of electrons with the N atom to form a double bond (see Figure 2 Step 5 above).
E. To finish off a polyatomic ion, place square brackets around the Lewis structure, and show the charge in the upper right-hand corner outside the brackets. The complete Lewis structure for the nitrite ion, NO2-, is shown in Figure 2 Step 6 on the previous page.
F. Finally, confirm that electrons were neither created nor destroyed – that there are in fact 14 electrons, no more or less, in the final Lewis structure.
Resonance Structures
Note that the double bond on the nitrite ion can be made with either oxygen atom, so two different Lewis structures can be drawn for the nitrite ion, as shown below:
H C Cl
H
H
H C Cl
H
H
H C Cl
H
H
O N.. O..
.. .. O^ N..^ O....
O N O.. ..
O N.. O....
O N O
Valence-shell electron-pair repulsion (VSEPR) model
Lewis structures show the two-dimensional distribution of atoms and electrons. The molecular geometry, or three- dimensional shape of a molecule or polyatomic ion, can be determined using valence-shell electron-pair repulsion (abbreviated VSEPR and pronounced “VES-per”) theory. The basic principle of VSEPR is that valence electrons around a central atom stay as far apart as possible to minimize the repulsions.
General Rules for the VSEPR Model Consider double and triple bonds to be single electron “groups” that will repel other bonds in the molecule. If two or more resonance structures can be drawn for a molecule, VSEPR model can be applied to any one of them.
Guidelines for Applying the VSEPR Model
- Draw the Lewis structure
- Count the number of outer atoms and lone pairs around the central atom.
- Match the bonded pairs and lone pair electrons to the molecular geometry, name and bond angle(s) for the molecule, using a reference table or your own knowledge.
- Remember that lone pair electrons occupy more space than bonded pairs of electrons, so a central atom’s lone pairs will compress other bond angles (e.g. <120, <109.5, etc.) around the central atom.
Example 1: Determine the molecular geometry for CH 3 Cl.
According to the Lewis structure, there are four outer atoms and no lone pairs on the central atom. The bonded electrons around the central carbon atom push apart from each other in 3 dimensions, resulting in a tetrahedral geometry which has
109.5 bond angles.
Thus, CH 3 Cl’s molecular geometry (or 3D shape) is tetrahedral, and its bond angles are 109.5 . Example 2: Determine the molecular geometry for the nitrite ion, NO 2 –.
NO 2 -^ has two resonance structures:
Both of these show two bonding groups of electrons around the central atom, as well as one lone pair on the central atom, for a total of three sets of electrons around the central atom. According to VSEPR, these three sets of electrons will push apart in three dimensions, producing a trigonal planar geometry of electrons. HOWEVER, because there are
only two outer atoms, the molecular geometry is bent with <120 bond angles. Recall also that lone pairs around the
central atom repel the bonding pairs of electrons to compress the bond angle. Thus, nitrite ion’s molecular geometry
is bent, and its bond angles are <120 .
H C Cl
..
....
H
H
O N.. O....
O N.. O....
Molecular Geometries
Image Credit: Chemistry, Open Stax. License: CC BY: Attribution
Polarity of Molecules
Electronegativity is the ability of an atom in a chemical bond to attract electrons to itself.
Trends for Electronegativity (EN) Fluorine (F) is the most electronegative element in the Periodic Table. In general, the closer an element is to fluorine, the more electronegative it is. The one exception is hydrogen (H) which has an electronegativity value between those of boron (B) and carbon (C).
Polar Covalent Bonds When two atoms share electrons in a covalent bond, they do not share equally if one atom is more electronegative than the other. The “tug of war” for the electrons results in the more electronegative atom pulling the electrons more strongly towards itself, so it becomes the negative (electron rich) end of the bond. Since electrons spend less time around the other atom, that atom becomes the positive (electron deficient) end of the bond. One atom is negative and the other is positive, so the bond has two poles (like the positive and negative ends of a magnet). This separation of charges is called a dipole , and bonds with a dipole are polar covalent bonds. In the example below, the very electronegative F pulls the bonding electrons away from the less electronegative H. This is shown using delta ( ) notation , where the more electronegative atom gets the –^ and the less electronegative atom gets the +.
H F
Electronegativity
How to draw molecules in 3D
Use the following notation to show 3D molecules. Solid bars show atoms coming out toward the viewer, while dashed bars show the molecule receding into the background. Regular straight lines show molecules in the plane of the paper.
Drawing Abbreviated Molecular Structures
Due to carbon’s 4 valence electrons, any carbon atom in a stable molecule will always make 4 covalent bonds. Since these bonds can be with other carbon atoms, carbon compounds can be very large to the point where drawing a lewis dot structure becomes cumbersome and messy due to the number of atoms involved. Therefore, the structures of carbon based compounds is typically abbreviated:
Lewis structure (left) and abbreviated structure (right) of Vitamin A
In these abbreviated structures, it is assumed that:
Every intersection of two lines at an angle represents a carbon atom Any carbon atom with fewer than 4 bonds has the appropriate number of hydrogen atoms bonded to it to give it 4 bonds Any atom (other than carbon or hydrogen) with fewer than 4 bonds has the appropriate number of lone pairs to complete the octet.
