Covalent bonds in a molecule and the overall charge of a molecule can be visualized with Lewis dot structures.
Calculate formal charges on atoms in a compound
Pictorial representations are often used to visualize electrons, as well as any bonding that may occur between atoms in a molecule. In particular, chemists use Lewis structures (also known as Lewis dot diagrams, electron dot diagrams, or electron structures) to represent covalent compounds. In these diagrams, valence electrons are shown as dots that sit around the atom; any bonds that the atoms share are represented by single, double, or triple lines.
Generally, most Lewis structures follow the octet rule; they will share electrons until they achieve 8 electrons in their outermost valence shell. However, there are exceptions to the octet rule, such as boron, which is stable with only 6 electrons in its valence shell. The elements hydrogen (H) and helium (He) follow the duet rule, which says their outermost valence shell is full with 2 electrons in it.
To draw a Lewis structure, the number of valence electrons on each atom in the compound must be determined. The total number of valence electrons in the entire compound is equal to the sum of the valence electrons of each atom in the compound. Non-valence electrons are not represented when drawing the Lewis structures.
Valence electrons are placed as lone pairs (two electrons) around each atom. Most atoms may have an incomplete octet of electrons. However, atoms can share electrons with each other to fulfill this octet requirement. A bond that shares two electrons is called a single bond and is signified by a straight, horizontal line.
If the octet rule is still not satisfied, atoms may form a double (4 shared electrons) or triple bond (6 shared electrons). Because the bonding pair is shared, the atom that had the lone pair still has an octet, and the other atom gains two or more electrons in its valence shell.
For example, CO2 is a neutral molecule with 16 total valence electrons. In the Lewis structure, carbon should be double-bonded to both oxygen atoms.
Lewis structures can also be drawn for ions. In these cases, the entire structure is placed in brackets, and the charge is written as a superscript on the upper right, outside of the bracket.
Although we know how many valence electrons are present in a compound, it is harder to determine around which atoms the electrons actually reside. To assist with this problem, chemists often calculate the formal charge of each atom. The formal charge is the electric charge an atom would have if all the electrons were shared equally.
The formal charge of an atom can be determined by the following formula:
In this formula, V represents the number of valence electrons of the atom in isolation, N is the number of non-bonding valence electrons, and B is the total number of electrons in covalent bonds with other atoms in the molecule.
For example, let's calculate the formal charge on an oxygen atom in a carbon dioxide (CO2) molecule:
FC = 6 valence electrons - (4 non-bonding valence electrons + 4/2 electrons in covalent bonds)
FC = 6 - 6 = 0
The oxygen atom in carbon dioxide has a formal charge of 0.
Sometimes multiple Lewis structures can be drawn to represent the same compound. These equivalent structures are known as resonance structures and involve the shifting of electrons and not of actual atoms. Depending on the compound, the shifting of electrons may cause a change in formal charges. Most often, Lewis structures are drawn so that the the formal charge of each atom is minimized.
Resonance structures depict possible electronic configurations; the actual configuration is a combination of the possible variations.
Describe how to draw resonance structures for compounds
Lewis dot structures can be drawn to visualize the electrons and bonds of a certain molecule. However, for some molecules not all the bonding possibilities cannot be represented by a single Lewis structure; these molecules have several contributing or "resonance" structures. In chemistry terms, resonance describes the fact that electrons are delocalized, or flow freely through the molecule, which allows multiple structures to be possible for a given molecule.
Each contributing resonance structure can be visualized by drawing a Lewis structure; however, it is important to note that each of these structures cannot actually be observed in nature. That is, the molecule does not actually go back and forth between these configurations; rather, the true structure is an approximate intermediate between each of the structures. This intermediate has an overall lower energy than each of the possible configurations and is referred to as a resonance hybrid. It is important to note that the difference between each structure lies in the location of the electrons and not in the arrangement of the atoms.
For example, the nitrate ion, NO3-, has more than one valid Lewis structure. The structure contains two N-O single bonds and one N=O double bond. But the question then remains as to which oxygen should be involved in the double bond. Therefore, three valid resonance structures can be drawn. Double-ended arrows are used to indicate that the structures are chemically equivalent. Again, in reality, the electronic configuration does not change between the three structures; rather, it has one structure in which the extra electrons are distributed evenly. These fractional bonds are sometimes depicted by dashed arrows, which show that the electron density is spread out throughout the compound.
When you are drawing resonance structures, it is important to remember to shift only the electrons; the atoms must have the same position. Sometimes, resonance structures involve the placement of positive and negative charges on specific atoms. Because atoms with electric charges are not as stable as atoms without electric charges, these resonance structures will contribute less to the overall resonance structure than a structure with no charges.