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Lewis structures, also called electron-dot structures or electron-dot diagrams, are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently-bonded molecule, as well as coordination compounds. The Lewis Structure was named after G.N. Lewis, who introduced it in his 1916 article The Molecule and the Atom.
Lewis structures show each atom in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another (rarely, pairs of dots are used instead of lines). Excess electrons that form lone pairs are represented as pair of dots, and are placed next to the atoms on which they reside.
Individual Atoms
The Lewis structure for an individual atom is drawn by placing a dot around the atom for each valence electron available. There are four positions available for dots to be placed; most chemists draw them on the top, left, bottom, and right of the atom.
Each of the four positions must receive one electron (one dot) before any electrons can be paired with one another. (Helium is an exception; the two electrons are shown as a pair.)
Molecules
Choosing a central atom
Most simple covalent molecules have one central atom that is surrounded by the other atoms in the molecule. For compounds comprised of a single atom of one element and multiple atoms of a second element, the single atom is the central atom. For compounds containing two or more elements with only one atom in the compound, the central atom is the least electronegative single atom that is not hydrogen. For instance, in thionyl chloride (SOCl2), the sulfur atom is the central atom.
The initial Lewis structure should be drawn with each non-central atom connected by a single bond to the central atom.
Molecules without a central atom
In some cases, the central atom is difficult to determine, generally when there is no element with only one atom in the compound. In these cases, the connectivity of the structure--that is, the determination of which atoms are connected to which others--must be determined in some other way, either by trial and error, prior knowledge, or experimentation.
When the connectivity between the atoms is known or determined, the Lewis structure can be drawn in exactly the same fashion.
Counting electrons
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.
The Octet Rule
The octet rule states that most atoms are most stable when their valence shell contains eight electrons, either lone pairs or bonded pairs. For this purpose, each bond counts as two electrons for both atoms it connects. When drawing a Lewis structure, care should be taken never to place more than eight electrons around a single atom.
There are several exceptions to the octet rule. Hydrogen never has more than two electrons in its valence shell--and thus in a Lewis structure may only be bonded once, and may not have lone pairs. Beryllium may have four electrons in the valence shell, and boron and aluminum may have six. Conversely, nonmetals of the third period or higher may form "extended octets"--that is, they may contain more than eight electrons in their valence shell, generally by placing the extra electrons in a d orbital.
The 18-electron rule and the 32-electron rule
The 18-electron rule is operative on atoms from period 4, which have to achieve 18 electrons to fill their orbitals and achieve a noble gas configuration. Similarly from period 6, the atoms have to achieve 32 electrons to fill their orbitals.
Placing electrons
Once the total number of available electrons has been determined, electrons must be placed into the structure. They should be placed initially as lone pairs: one pair of dots for each pair of electrons available. Lone pairs should initially be placed on outer atoms (other than hydrogen) until each outer atom has eight electrons in bonding pairs and lone pairs; extra lone pairs may then be placed on the central atom. When in doubt, lone pairs should be placed on more electronegative atoms first.
Once all lone pairs are placed, atoms, especially the central atoms, may not have an octet of electrons. In this case, the atoms must form a double bond; a lone pair of electrons is moved to form a second bond between the two atoms. As the bonding pair is shared between the two atoms, the atom that originally had the lone pair still has an octet; the other atom now has two more electrons in its valence shell.
Extended Octets
Atoms of the third period or higher may form extended octets by placing extra electrons into the d orbitals of the third energy level. In this way, phosphorus may form five bonds, sulfur six, and the halogens as many as seven.
Ions
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.
Formal Charge
A discussion of formal charge is necessary to complete some more complicated Lewis structures. The formal charge of an atom is the charge that it would have if every bond were 100% covalent (non-polar). Formal charges are computed by using a set of rules and are useful for accounting for the electrons when writing a reaction mechanism, but they don't have any intrinsic physical meaning. They may also be used for qualitative comparisons between different resonance structures (see below) of the same molecule, and often have the same sign as the partial charge of the atom, but there are exceptions.
FC = # valence electrons - (½ × # bonding electrons + # lone pair electrons)
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.
- Example: Ammonium, NH4+ is a cationic species. Using the vertical groups of the atoms on the periodic table we determine that each hydrogen contributes 1 electron, the nitrogen contributes 5 electrons, and the charge of +1 means that 1 electron is absent. The final total is 8 total electrons (1 × 4 + 5 − 1). Drawing the Lewis structure gives an sp3 (4 bonds) hybridized nitrogen atom surrounded by hydrogen. There are no lone pairs of electrons left. Thus, using our definition of formal charge, hydrogen has a formal charge of zero (1-(½ × 2 - 0)) and nitrogen has a formal charge of +1 (5−(½ × 8 - 0)). After adding up all the formal charges throughout the molecule the result is a total formal charge of +1, consistent with the charge of the molecule given in the first place.
Note: The total formal charge in a molecule should be as close to zero as possible, with as few charges on the molecule as possible
- Example: CO2 is a neutral molecule with 16 total electrons. There are three different ways to draw the Lewis structure
- Carbon single bonded to both oxygen atoms (carbon = +2, oxygens = -1 each, total formal charge = 0)
- Carbon single bonded to one oxygen and double bonded to another (carbon = +1, oxygendouble = 0, oxygensingle = −1, total formal charge = 0)
- Carbon double bonded to both oxygen atoms (carbon = 0, oxygens = 0, total formal charge =0)
Even though all three structures gave us a total charge of zero, the final structure is the superior one because there are no charges in the molecule at all.
Resonance
For some molecules and ions, it is difficult to determine which lone pairs should be moved to form double or triple bonds. 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 an average of these states.
The nitrate ion (NO3-), for instance, must form a 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.
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.
Resonance Hybridization
In Lewis resonance structures, the structure is written such that it appears the molecule may switch between multiple forms. However, the molecule itself exists as a hybrid of the forms.
In the case of the nitrate ion, there are two single bonds and one double bond in each resonance form. When the nitrate ion is examined, however, each bond appears as though it has a bond order of 1.333--that is, each has properties as if it were composed of one and one-third total bonds. The length and bond energy of each is somewhere between that of a single bond and a double bond.
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.
Example: The Lewis structure of the nitrite ion
The formula of the nitrite ion is NO2−.
Step one: Choose a central atom. There is only one nitrogen atom, and it is the least electronegative atom, so it is the central atom by multiple criteria.
Step two: Count valence electrons. Nitrogen has 5 valence electrons; each oxygen has 6, for a total of (6 × 2) + 5 = 17. The ion has a charge of −1, which indicates an extra electron, so the total number of electrons is 18.
Step three: Place ion pairs. Each oxygen must be bonded to the nitrogen, which uses four electrons — two in each bonds. The 14 remaining electrons should initially be placed as 7 lone pairs. Each oxygen may take a maximum of 3 lone pairs, giving each oxygen 8 electrons including the bonding pair. The seventh lone pair must be placed on the nitrogen atom.
Step four: Satisfy the octet rule. Both oxygen atoms currently have 8 electrons assigned to them. The nitrogen atom has only 6 electrons assigned to it. One of the lone pairs on an oxygen atom must form a double bond, but either atom will work equally well. We therefore must have a resonance structure.
Step five: Tie up loose ends. Two Lewis structures must be drawn: one with each oxygen atom double-bonded to the nitrogen atom. The second oxygen atom in each structure will be single-bonded to the nitrogen atom. Place brackets around each structure, and add the charge (−) to the upper right outside the brackets. Draw a double-headed arrow between the two resonance forms.
Alternative formats
Chemical structures may be written in more compacted forms, particularly when showing organic molecules. In condensed structural formulas many or even all of the covalent bonds may be left out and replaced with subscripts indicating the number of identical groups attached to a particular atom. Another, strictly diagramic, alternate form is the "bond-line formula" or "carbon skeleton diagram". In bond-line formulas, the only atoms specified explicitly are those that are neither carbon nor hydrogen bound to carbon.
External links
- Think Quest run by Thinkquest offers more practice and examples with Lewis structures.
References
- Lewis, G. N. The Atom and the Molecule. J. Am. Chem. Soc. 1916, 38, 762-785. *
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"Lewis structure"