BNG_Lewis Symbols Lesson

Lewis Symbols

When studying bonding, we know that the valence electrons are most important. To keep track of valence electrons, Lewis symbols  are used. They were developed by a famous American scientist, G.N. Lewis (1875-1946) as a bookkeeping method for keeping track of valence electrons.  

To draw the Lewis symbol for an element, write its chemical symbol surrounded by dots to represent each valence electron. Elements that are in the same group on the periodic table have the same configuration of dots in their Lewis symbol. For example, elements in group 1 (the alkali metals) all have one valence electron. This means that they will have one dot in their Lewis symbol like this, X·

The Lewis symbol for the representative elements in period 2 are below.

Group 1 Li Group 2 Be Group 13 B Group 14 C Group 15 N Group 16 0 Group 17 F Group 18: Ne

Notice that in group 15, when there are 5 valence electrons, the electrons begin to pair up. One side has a pair of electrons, while the other 3 sides have a single electron. When drawing the Lewis symbol, it does not matter which side has the pair of electrons and which sides have the lone electrons.

Lewis Structures

It is easy to see how ionic bonds are formed since there is a complete transfer of electrons. When electrons are shared to form covalent bonds, it isn't always so clear. In order to determine exactly how the electrons are shared, we will use Lewis symbols to draw Lewis structures. Lewis structures  are used to show how the valence electrons are arranged in a molecule. Watch the following to learn how to draw Lewis structures.

Practice drawing Lewis structures on your own by using the interactive below. You can check your structures as you work. For the double and triple bond examples, you should do those on your own paper (or computer program) and then click on each molecule to check your answers. You will click on the molecule twice to see how the double bonds are formed and three times to see how the triple bonds are formed.

doubletriplebonds.pngLet's remember that these structures can be simplified a bit by replacing the bonded electrons with a dash. For example, in the H2 molecule, H:H becomes H-H. This is an example of a  single bond. One pair of electrons is shared between two atoms. As we have already seen in our examples, some molecules form  double or triple bonds. This happens when there are not enough valence electrons available for the octet rule to be followed through only single bonds. CO2 has two double bond, and N2 has one triple bond as shown to the right.

Bond Order, Bond Length, and Bond Energy

Once a Lewis structure of a molecule is drawn, something called bond order can be determined. Bond order is the number of bonding pairs of electrons between two atoms. So, a single bond has a bond order of one, a double bond has a bond order of 2, etc. If you are asked about a molecule's bond order, you will need to draw that Lewis structure in order to answer the question.

*To determine the bond order of a molecule with more than two elements, you will follow these steps:

  1. Draw the Lewis structure
  2. Add up the total number of bonds (a double bond counts as two bonds ...)
  3. Add up the number of bond clouds (here a double bond or triple bond counts as just one cloud)
  4. Divide the number of bonds by the number of bonding clouds

 

Bond Length and Bond Order

The average distance between the two nuclei is called bond length. The smaller the bond length, the closer the nuclei are to one another. Consequently, the smaller the bond length, the stronger the bond. Bond order is related to bond length inversely. That is, as the bond order increases, the bond length decreases.    

Bond length and bond order both tell us something about the strength of the covalent bond.  

  • The shorter the bond length, the stronger the bond.
  • The higher the bond order, the stronger the bond.

So, short bond length and high bond order both indicate that the atoms are held together tighter than those with longer bond length and smaller bond order.  

Bond Strength and Bond Energy

Bond strength can be quantified by something called bond energy. Since a covalent bond holds atoms together, it takes energy to separate them. The amount of energy required to break a bond is called bond dissociation energy   or simply bond energy.

The bond order is related to the bond energy directly. As the bond order increases, the bond energy also increases.

BrokenPencil.jpeg"When a bond is strong, there is a higher bond energy because it takes more energy to break a strong bond. This correlates with Bond Order and Bond Length. When the Bond Order is higher, Bond Length is shorter, and the shorter the Bond Length means a greater the Bond Energy because of increased electric attraction. Think about it this way: it is easy to snap a pencil, but if you keep snapping the pencil it gets harder each time since the length of the pencil decreases. A higher bond energy (or a higher bond order or shorter bond length) means that a bond is less likely to break apart. In other words, it is more stable than a molecule with a lower bond energy. With Lewis Structures then, the structure with the higher bond energy is  more likely to occur." 

We will do calculations with bond energy in a later unit.

Limitations to Lewis Structure

As with any model, there are limitations to the use of the Lewis structure model, particularly in cases with an odd number of valence electrons. You should recognize that Lewis diagrams have exceptions:

Expanded Octet

This is when an element forms bonds that give them more than 8 valence electrons. This can only happen if the central element is on the 3rd energy level or higher. These elements often have an expanded valence because of their larger radii (when compared to the second-row elements) and the availability of empty d-orbitals in the valence shell. Remember that d-orbitals do not exist on the 1st or 2nd energy level (or shell). These atoms tend to extend their octet when the bonded atoms are highly electronegative, such as O, F,or Cl.

Fewer than eight valence electrons

Electron-deficient compounds are compounds in which an element has an incomplete octet. Some elements, such as H, Be, and B, often have fewer than eight electrons in Lewis structures. Hydrogen has a single valence electron in a 1s orbital and therefore can only form one bond. Beryllium (two valence electrons) and boron (three valence electrons) often accommodate only four or six electrons, respectively, in Lewis structures. Examples are BeH2, and BF3

Odd-electron compounds

A few stable compounds contain an odd number of valence electrons and thus cannot obey the octet rule. NO, NO2, and are some examples ClO2.

While you do need to be able to draw the Lewis structures for such exceptions, learning how to defend Lewis models based on assumptions about the limitations of the models is beyond the scope of this course and the AP Exam. In other words, you do not have to explain why an element forms an expanded octet, only that it is possible to do so because of their larger radii (when compared to the second-row elements) and the availability of empty d-orbitals in the valence shell.

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