Lewis structuresalso called Lewis dot formulas, Lewis dot structures, electron dot structures, or Lewis electron dot structures (LEDs)are diagrams that show the bonding between atoms of a molecule, as well as the lone pairs of electrons that may exist in the molecule.[1] [2] A Lewis structure can be drawn for any covalently bonded molecule, as well as coordination compounds. The Lewis structure was named after Gilbert N. Lewis, who introduced it in his 1916 article The Atom and the Molecule. Lewis structures extend the concept of the electron dot diagram by adding lines between atoms to represent shared pairs in a chemical bond.
Lewis structures show each atom and its position in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another (pairs of dots can be used instead of lines). Excess electrons that form lone pairs are represented as pairs of dots, and are placed next to the atoms.
Although main group elements of the second period and beyond usually react by gaining, losing, or sharing electrons until they have achieved a valence shell electron configuration with a full octet of (8) electrons, hydrogen (H) can only form bonds which share just two electrons.
See main article: article and Electron counting. The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons on each individual atom. Non-valence electrons are not represented in Lewis structures.
Once the total number of valence electrons has been determined, they are placed into the structure according to these steps:
Lewis structures for polyatomic ions may be drawn by the same method. When counting electrons, negative ions should have extra electrons placed in their Lewis structures; positive ions should have fewer electrons than an uncharged molecule. When the Lewis structure of an ion is written, the entire structure is placed in brackets, and the charge is written as a superscript on the upper right, outside the brackets.
A simpler method has been proposed for constructing Lewis structures, eliminating the need for electron counting: the atoms are drawn showing the valence electrons; bonds are then formed by pairing up valence electrons of the atoms involved in the bond-making process, and anions and cations are formed by adding or removing electrons to/from the appropriate atoms.
A trick is to count up valence electrons, then count up the number of electrons needed to complete the octet rule (or with hydrogen just 2 electrons), then take the difference of these two numbers. The answer is the number of electrons that make up the bonds. The rest of the electrons just go to fill all the other atoms' octets.
Another simple and general procedure to write Lewis structures and resonance forms has been proposed.
There is a way to construct Lewis Structures reliably via the use of a table similar to the one below:
| Molecular Formula (CCl4) | Octet electron (Oe) | Total Valence electrons (TVe) | |
|---|---|---|---|
| C | 8 | 4 | |
| Cl | 8 | 7 | |
| Cl | 8 | 7 | |
| Cl | 8 | 7 | |
| Cl | 8 | 7 | |
| total: | 40 | 32 |
Bond number is calculated as follows: (Oe - TVe) ÷ 2
For the example of CCl4, bond number would be: (40 - 32) ÷ 2 = 8 ÷2 = 4. Therefore, CCl4 has 4 bonds.
Lone pairs are not directly calculated, instead one calculates the number of remaining electrons. This is done as follows: TVe - 2(number of bonds)
For the example of CCl4, the number of remaining electrons would be: 32 - 2(4) = 32 - 8 = 24. Therefore, CCl4 has 24 remaining electrons, which means it has 12 lone pairs.
This system works in nearly all cases, however there are 3 instances where it will not work. These exceptions are outlined in the table below.
| The Exception | How it Breaks the System | How to Fix the Lewis Structure | |
|---|---|---|---|
| Free Radicals (molecules with unpaired valence electrons) | Sum of TVe will be an odd number. Bond number will not be a whole number. | Round calculated bond number down to the nearest whole number. (e.g. 4.5 bonds would round down to 4 bonds) | |
| Valence Shell Deficiency | Does not break the system, must instead memorize when it occurs. | BeX2, BX3, and AlX3. "X" represents Hydrogen or Halogens. When Be is bonded with 2 other atoms, or when B and Al are bonded with 3 other atoms, they do not form full valence shells. Assume single bonds and use the actual bond number to calculate lone pairs. | |
| Expanded Octet (only occurs for elements in Groups 3-8) | Bond calculation will provide too few bonds for the number of atoms in the molecule. | Assume single bonds, use the minimum number of bonds necessary to create the molecule. |
See main article: article and Formal charge. In terms of Lewis structures, formal charge is used in the description, comparison, and assessment of likely topological and resonance structures[3] by determining the apparent electronic charge of each atom within, based upon its electron dot structure, assuming exclusive covalency or non-polar bonding. It has uses in determining possible electron re-configuration when referring to reaction mechanisms, and often results in the same sign as the partial charge of the atom, with exceptions. In general, the formal charge of an atom can be calculated using the following formula, assuming non-standard definitions for the markup used:
Cf=Nv-Ue-
| Bn | |
| 2 |
Cf
Nv
Ue
Bn
The formal charge of an atom is computed as the difference between the number of valence electrons that a neutral atom would have and the number of electrons that belong to it in the Lewis structure. Electrons in covalent bonds are split equally between the atoms involved in the bond. The total of the formal charges on an ion should be equal to the charge on the ion, and the total of the formal charges on a neutral molecule should be equal to zero.
See main article: article and Resonance structure. For some molecules and ions, it is difficult to determine which lone pairs should be moved to form double or triple bonds, and two or more different resonance structures may be written for the same molecule or ion. In such cases it is usual to write all of them with two-way arrows in between . This is sometimes the case when multiple atoms of the same type surround the central atom, and is especially common for polyatomic ions.
When this situation occurs, the molecule's Lewis structure is said to be a resonance structure, and the molecule exists as a resonance hybrid. Each of the different possibilities is superimposed on the others, and the molecule is considered to have a Lewis structure equivalent to some combination of these states.
The nitrate ion, for instance, must form a double bond between nitrogen and one of the oxygens to satisfy the octet rule for nitrogen. However, because the molecule is symmetrical, it does not matter which of the oxygens forms the double bond. In this case, there are three possible resonance structures. Expressing resonance when drawing Lewis structures may be done either by drawing each of the possible resonance forms and placing double-headed arrows between them or by using dashed lines to represent the partial bonds (although the latter is a good representation of the resonance hybrid which is not, formally speaking, a Lewis structure).
When comparing resonance structures for the same molecule, usually those with the fewest formal charges contribute more to the overall resonance hybrid. When formal charges are necessary, resonance structures that have negative charges on the more electronegative elements and positive charges on the less electronegative elements are favored.
Single bonds can also be moved in the same way to create resonance structures for hypervalent molecules such as sulfur hexafluoride, which is the correct description according to quantum chemical calculations instead of the common expanded octet model.
The resonance structure should not be interpreted to indicate that the molecule switches between forms, but that the molecule acts as the average of multiple forms.
The formula of the nitrite ion is .
Chemical structures may be written in more compact forms, particularly when showing organic molecules. In condensed structural formulas, many or even all of the covalent bonds may be left out, with subscripts indicating the number of identical groups attached to a particular atom.Another shorthand structural diagram is the skeletal formula (also known as a bond-line formula or carbon skeleton diagram). In a skeletal formula, carbon atoms are not signified by the symbol C but by the vertices of the lines. Hydrogen atoms bonded to carbon are not shown—they can be inferred by counting the number of bonds to a particular carbon atom—each carbon is assumed to have four bonds in total, so any bonds not shown are, by implication, to hydrogen atoms.
Other diagrams may be more complex than Lewis structures, showing bonds in 3D using various forms such as space-filling diagrams.
Despite their simplicity and development in the early twentieth century, when understanding of chemical bonding was still rudimentary, Lewis structures capture many of the key features of the electronic structure of a range of molecular systems, including those of relevance to chemical reactivity. Thus, they continue to enjoy widespread use by chemists and chemistry educators. This is especially true in the field of organic chemistry, where the traditional valence-bond model of bonding still dominates, and mechanisms are often understood in terms of curve-arrow notation superimposed upon skeletal formulae, which are shorthand versions of Lewis structures. Due to the greater variety of bonding schemes encountered in inorganic and organometallic chemistry, many of the molecules encountered require the use of fully delocalized molecular orbitals to adequately describe their bonding, making Lewis structures comparatively less important (although they are still common).
There are simple and archetypal molecular systems for which a Lewis description, at least in unmodified form, is misleading or inaccurate. Notably, the naive drawing of Lewis structures for molecules known experimentally to contain unpaired electrons (e.g., O2, NO, and ClO2) leads to incorrect inferences of bond orders, bond lengths, and/or magnetic properties. A simple Lewis model also does not account for the phenomenon of aromaticity. For instance, Lewis structures do not offer an explanation for why cyclic C6H6 (benzene) experiences special stabilization beyond normal delocalization effects, while C4H4 (cyclobutadiene) actually experiences a special destabilization. Molecular orbital theory provides the most straightforward explanation for these phenomena.