MICSP - Lewis Diagrams (Lesson)

Lewis Diagrams

Finding ways to visualize compounds, since we cannot see molecules with our eyes has given rise to several ways of representing matter. One of the most helpful is a Lewis Dot Diagram named after Gilbert Lewis.

These models best represent shared electrons in covalent bonds, but can be used for atoms, ions, and even salts.

Lewis Dot Structures:

• Devised by G.N. Lewis for representing valence electrons around atoms.
• Uses the element symbol with dots to represent valence electrons.
• Also used to show how electrons are shared in covalent bonds.

Atoms:

• Draw only outer electrons that fill in the groups 1,2,13-18.
• Atoms with d level electrons will not be represented with Lewis structures.
• Only valence electrons: Maximum of 8 electrons.
• In the Lewis symbol for an atom, the chemical symbol of the element (as found on the periodic table) is written, and the valence electrons are represented as dots surrounding it. Only the electrons in the valence level are shown using this notation. Think of the symbol enclosed in by a square - you may only place 1 or 2 electrons on each side of the square - one side represents the outermost s energy sublevel, and the other three sides the individual p valence orbitals.

For example, the Lewis symbol of carbon depicts a “C" surrounded by 4 valence electrons because carbon has an electron configuration of 1s²2s²2p².

 

Simple Ions

Simple Ions are atoms that gain or lose electrons to reach a noble gas configuration. For example, the sodium ion would not have any dots and a + sign. The symbol would then be enclosed in brackets representing an empty outer valence level. 

The chloride ion would show eight dots, with the negative sign to show the filled valence level. Again, enclosed with brackets showing the full outer valence shell.

Simple Molecular Compounds with Covalent Bonds

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. A covalent bond is a pair of electrons (usually one from each atom) shared between two nuclei.

• Since shared electrons contribute to the valence electron count of both atoms bonding, the number of possible bonds of any one atom has a maximum of four: as there are two electrons minimum in a bond and a total of 8 valence electrons possible. (8/2 = 4).

Example: The unpaired e- of a chlorine atom often pairs with an unpaired e- of another atom to form one covalent bond. This gives Cl an octet (2 e- from the two‑electron covalent bond, and 6 from its three lone pairs). This is why Cl gas exists as Cl₂ molecules.

Please watch the following presentation to learn how to draw Lewis structures. Be sure your volume is turned up!

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.

Multiple Bonds Consist of Two or More Pairs of Electrons

The bond produced by the sharing of one pair of electrons between two atoms is called a single bond. So far, these have been the only kind we've discussed. There are, however, many molecules in which more than a single pair of electrons are shared between two atoms. For one, just exhale... yes, the carbon dioxide produced in respiration has carbon sharing two of its four valence electrons with each of the two oxygen molecules.

Sometimes, three pairs of electrons are shared between two atoms. The most abundant gas in our atmosphere, nitrogen, occurs in the form of diatomic molecules of N₂.

The bond energy is greater in a double bond than in a single bond and a triple bond has the highest of the three types. The bond length is shorter for a multiple bond than for a single bond, indicating that the nuclei of the atoms sharing electrons are closest in a triple bond.

Bond Order is a term used to note single or multiple bonds with a number 1 for single, 2 for double, etc.

Carbon Chains

With four valence electrons, carbon often forms long chains with other carbons.

• In a correct structure, counting shared electrons with both atoms, H always winds up with 2 outer electrons and every other atom with 8. Both atoms in the bond must come from groups 4-7, including H.  Carbon tends to form chains of atoms with itself. 

• Propane:  C₃H₈ also written as CH₃CH₂CH₃. Two Carbons bonded to each other surrounded by H’s. 

• Propanol: For compounds like C₃H₇OH below, the carbons link to each other, and the other atoms are attached to the chain.

• Formulas are also written for carbon chains to show a structural hint on how to arrange the groups – C₃H₇OH could also be written CH₃CH₂CH₂OH to show that three H’s are on the first, two H’s on the second carbon and 2 H’s and the OH are attached to the third carbon.

Limitations to the Lewis Structure Model

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 3^rd 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 1^st or 2^nd energy level (or shell). These atoms tend to expand their octet when bonded to atoms that 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 as well as ClO₂. Structures like these are best determined using Formal Charge calculations that you will find in the next lessons of this module.

You Try It!

Write Lewis symbols for the following atoms. Click on the plus sign to reveal the answer to check your work!

Write Lewis structures for the following compounds. Click on the plus sign to reveal the answer to check your work!

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