A Lewis electron dot diagram (or electron dot diagram or a Lewis diagram or a Lewis structure) is a representation of the valence electrons of an atom that uses dots around the symbol of the element.
Electrons exist outside of an atom‘s nucleus and are found in principal energy levels that contain only up to a specific number of electrons. The outermost principal energy level that contains electrons is called the valence level and contains valence electrons. Lewis symbols and formulas use dots to visually represent the valence electrons of an atom and can be used to represent covalent bonds as shared pairs of electrons between atoms. Lewis symbols do not visualize the electrons in the inner principal energy levels.
The number of dots equals the number of valence electrons in the atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots per side. (It does not matter what order the positions are used.) When drawing more complex covalent compounds we need to also consider the formal charge and resonance structures of each compound.
For example, the Lewis electron dot diagram for calcium is simply:
Lewis symbols can also be used to illustrate the formation of cations from atoms, as shown here for sodium and calcium:
Likewise, they can be used to show the formation of anions from atoms, as shown below for chlorine and sulfur:
When considering how atoms come together to form covalent bonds we need to consider Lewis diagrams which can illustrate how atoms come together to share pairs of electrons. Two H atoms can come together and share each of their electrons to create a covalent bond. The shared pair of electrons can be thought of as belonging to either atom and thus each atom now has two electrons in its valence level.
In many atoms, not all of the electron pairs comprising the octet are shared between atoms. These unshared, non-bonding electrons are called lone pairs of electrons. Although lone pairs are not directly involved in bond formation, they should always be shown in Lewis structures.
Lewis structures can be drawn in 4 steps:
- Write a structural diagram of the molecule to clearly show which atom is connected to which (although many possibilities exist, we usually pick the element with the most number of possible bonds to be the central atom).
- Draw Lewis symbols of the individual atoms in the molecule.
- Bring the atoms together in a way that places eight electrons around each atom (or two electrons for H, hydrogen) wherever possible.
- Each pair of shared electrons is a covalent bond that can be represented by a dash.
Lewis structures can also be drawn for polyatomic ions. The Lewis structure of an ion is placed in brackets and its charge is written as a superscript outside of the brackets, on the upper right. Ions are treated almost the same way as a molecule with no charge. However, the number of electrons must be adjusted to account for the net electric charge of the ion. When counting electrons, negative ions should have extra electrons placed in their Lewis structures, while positive ions should have fewer electrons than an uncharged molecule.
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 can be calculated using the formula:
For example, the formal charge of an oxygen atom in carbon dioxide is:
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 formal charge of each atom is minimized. Resonance structures have the same number of electrons and therefore have the same overall charge. Resonance structures differ only in the arrangement of electrons; the atoms keep the same connectivity and arrangement.
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.
G. N. Lewis proposed a generalized definition of acid-base behavior in which acids and bases are identified by their ability to accept or to donate a pair of electrons and form a coordinate covalent bond.
A coordinate covalent bond (or dative bond) occurs when one of the atoms in the bond provides both bonding electrons. For example, a coordinate covalent bond occurs when a water molecule combines with a hydrogen ion to form a hydronium ion. A coordinate covalent bond also results when an ammonia molecule combines with a hydrogen ion to form an ammonium ion. Both of these equations are shown here.
A Lewis acid is any species (molecule or ion) that can accept a pair of electrons, and a Lewis base is any species (molecule or ion) that can donate a pair of electrons. A Lewis acid-base reaction occurs when a base donates a pair of electrons to an acid. A Lewis acid-base adduct, a compound that contains a coordinate covalent bond between the Lewis acid and the Lewis base, is formed. The following equations illustrate the general application of the Lewis concept.
• Lewis symbols are diagrams that show the number of valence electrons of a particular element with dots that represent lone pairs.
• Lewis structures incorporate an atom’s formal charge, which is the charge on an atom in a molecule, assuming that electrons in a chemical bond are shared equally between atoms.
• Lewis dot diagrams are often employed to visualize the covalent bonding between atoms in a compound. However, when multiple equally valid structures can be drawn, these structures are called resonance structures. Resonance structures have the same number of electrons and therefore have the same overall charge.
• A Lewis acid is any species (molecule or ion) that can accept a pair of electrons, and a Lewis base is any species (molecule or ion) that can donate a pair of electrons.
Valence level: The outermost principal energy level, which is the level furthest away from the nucleus that still contains electrons.
Valence electrons: The electrons of atoms that participate in the formation of chemical bonds.
Lewis symbols: Symbols of the elements with their number of valence electrons represented as dots
Electronegativity: The tendency of an atom or molecule to attract electrons and form bonds.
Covalent bond: A type of chemical bond where two atoms are connected to each other by the sharing of two or more electrons.
Cation: A positively charged ion.
Anion: A negatively charged ion.
Polyatomic ion: A charged species composed of two or more atoms covalently bonded, or of a metal complex that acts as a single unit in acid-base chemistry or in the formation of salts. Also known as a molecular ion.
Formal charge: The charge assigned to an atom in a molecule, assuming that electrons in a chemical bond are shared equally between atoms. This helps determine which of a few Lewis structures is most correct.
Resonance structure: A molecule or polyatomic ion that has multiple Lewis structures because bonding can be shown in multiple ways.
Coordinate covalent bond: A coordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.
Lewis acid: Any substance, such as the H+ ion, that can accept a pair of nonbonding electrons; an electron-pair acceptor.
Lewis base: Any substance, such as the OH- ion, that can donate a pair of nonbonding electrons; an electron-pair donor.
Lewis acid-base adduct: A compound that contains a coordinate covalent bond between the Lewis acid and the Lewis base.