7.4 Lewis Symbols and Structures

What is Lewis Electron Dot Structure?

Lewis Dot Structure is a simple presentation of the overlapping orbitals forming sigma and pi bond. The valence electrons contributing to the bonding (single, double and triple bonds) are presented as dots. Lewis Dot Structure illustrates the distribution of the valence electrons and the octet rule satisfaction to the atoms contributing to the bonding.

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.

Example Phosphorous and Argon Dot Structures:

Phosphorous has 5 valence electrons and Argon has 8 valence electrons.

Figure 7.46 Lewis Symbol of atoms

Ref: commons.wikimedia.org/

We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons:

Figure 7.47 Lewis Symbol of Ca

Ref: commons.wikimedia.org/

The following Table shows the Lewis symbols for the elements of the third period of the periodic table. Electron dots are typically arranged in four pairs located on the four “sides” of the atomic symbol.

Reference: https://www.youtube.com/watch?v=cIuXl7o6mAw&vl=en

How to Draw Lewis Electron Dot Structure?

A Lewis electron dot structure describes the bonding atoms, the number of bonds in the molecule, and the lone pairs left in the bonding atoms.

The general steps are given below:

1. First, the total number of valence electrons present in the molecule is calculated by adding the individual valencies of each atom.
2. If the molecule is an anion, extra electrons (number of electrons added = the magnitude of negative charge) are added to the Lewis dot structure.
3. When cationic molecules are considered, electrons are subtracted from the total count in order to make up for the positive charge.
4. The least electronegative atom is made the central atom of the molecule or ion.
5. The atoms are now connected via single bonds.
6. Now, the lone pairs of electrons are assigned to each atom belonging to the molecule. Commonly, the lone pairs are assigned to the most electronegative atoms first.
7. Once the lone pairs are assigned, if every atom does not have an octet configuration, a double or triple bond must be drawn to satisfy the octet valency of each atom.
8. If required, a lone pair can be converted into a bond pair in order to satisfy the octet rule for two atoms.

It is important to note that only the valence electrons are considered while drawing Lewis dot structures and the electrons that do not belong to the outermost shell are ignored. Examples are shown below.

Figure 7.48 Lewis Structure of covalent molecules

Ref: commons.wikimedia.org/

After reviewing all these rules, the Lewis Dot Structure of NO₃ ^–will be written as shown below:

Reference: https://www.youtube.com/watch?v=toXtRCEUBKU

Lewis Structure Examples

The Lewis electron dot structures of a few molecules are illustrated in this subsection.

The video below illustrates the Lewis Dot structure of some organic and inorganic compounds:

1. Lewis Structure of CO₂

• The central atom of this molecule is carbon.
• Oxygen contains 6 valence electrons which form 2 lone pairs. Since it is bonded to only one carbon atom, it must form a double bond.
• Carbon contains four valence electrons, resulting in zero lone pairs. Therefore, it is doubly bonded to each oxygen atom.

Figure 7.49 Lewis Dot Structure of CO₂

Ref: commons.wikimedia.org/

2. Lewis Structure of O₂

• An atom of oxygen contains 6 electrons in the valence shell.
• Four of thevalence electrons exist in lone pairs, implying that the oxygen atom must participate in two single bonds or one double bond in order to attain an octet configuration.
• Since there are only two oxygen atoms in an O₂ molecule, the atoms form a double bond resulting in the following Lewis electron dot structure.

Figure 7.50 Lewis Dot Structure of O₂

Ref: commons.wikimedia.org/

3. Lewis Structure of CO (Carbon Monoxide)

• A carbon monoxide molecule consists of one carbon atom and one oxygen atom.
• The carbon atom requires four electrons to obtain octet configuration whereas the oxygen atom requires two.
• Therefore, the valency is satisfied via the donation of a lone pair of electrons for bonding by the oxygen atom.
• The resulting Lewis electron dot structure displays atriple bond connecting a carbon and an oxygen atom, each holding a lone pair of electrons.

Figure 7.51 Lewis Dot Structure of CO

Ref: commons.wikimedia.org/

The Octet Rule Within Covalent Bonds

The octet rule is a chemical rule of thumb that reflects the observation that main group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas.

The octet rule dictates that atoms are most stable when their valence shells are filled with eight electrons. It is based on the observation that the atoms of the main group elements tend to participate in chemical bonding in such a way that each atom of the resulting molecule has eight electrons in the valence shell. The octet rule is only applicable to the main group elements.

The molecules of the halogens, oxygen, nitrogen, and carbon are known to obey the octet rule. In general, the elements that obey this rule include the s-block elements and the p-block elements (with the exception of hydrogen, helium, and lithium).

The octet rule can be observed in the bonding between the carbon and oxygen atoms in a carbon dioxide molecule, as illustrated via a Lewis dot structure below.

                                                Figure 7.52 The Octet Rule

The shared electrons fulfil the valency requirements of both the bonded atoms. Thus, it can be noted that both the oxygen atoms and the carbon atom have an octet configuration in the CO₂ molecule.

Upon observing that the noble gases were chemically inert, the electronic theory of valency was proposed by the German physicist Walther Kossel and the American chemist Gilbert Lewis. It was based on the tendency of atoms to assume the most stable state possible.

Illustrated below here for carbon in CCl₄ (carbon tetrachloride) and silicon in SiH₄ (silane). Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:

Figure 7.53 Lewis Dot structures

Ref: commons.wikimedia.org/

Lewis Structures: Single, Double & Triple Bonds

A Lewis dot structure can be represented by either two dots or by a line between two atoms when there is a chemical bond – two lines for double bond and three lines for a triple bond. … Connect the atoms by electron pair bonds so that each atom has a full octet.

Reference: https://www.youtube.com/watch?v=P5v7m-3WArw

So, in definition, Single Bond: “A covalent bond between two atoms that is formed by one bonding pair of electrons.” Double Bond: “A multiple covalent bond between two atoms that is formed by two bonding electron pairs. Triple Bond: “A multiple covalent bond between two atoms that is formed by three bonding pairs of electrons.

Multiple Covalent Bond: “more than one bond between atoms, formed by the sharing of more than one electron from each atom. Can be either a double bond or triple bond.” 

Examples:

A double bond

Figure 7.54 Double bond

Ref: commons.wikimedia.org

A triple bond

Figure 7.55 Lewis Dot structures

Ref: commons.wikimedia.org

Reference: https://www.youtube.com/watch?v=aG3di-2Pxmo

Reference: https://www.youtube.com/watch?v=SkyrRxtGo-A

Writing Lewis Structures with the Octet Rule

For individual atoms, the Lewis structure is drawn by placing a dot around the atom for each valence electron available. Octet rule states that in forming compounds, atoms gain, lose or share electrons to give a stable electron configuration characterized by eight valence electrons.

For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:

Figure 7.56 Satisfying Octet Rule using Lewis Dot structures

Ref: commons.wikimedia.org

For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:

1. Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.
2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).
3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.
4. Place all remaining electrons on the central atom.
5. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

Reference: https://www.youtube.com/watch?v=gTSXzAYcP4g

Reference: https://www.youtube.com/watch?v=96L6_IwyHRM

Bonding and nonbonding electrons: lone pairs

Figure 7.57 Bonding and nonbonding electrons

In many molecules, not all of the electron pairs comprising the octet are shared between atoms. The unshared electron pairs are often called lone pairs. Although lone pairs are not directly involved in bond formation, they should always be shown in Lewis formulas; we will see later that they have an important role to play in determining the shape of the molecule.

Deviation from the Octet Rule

Expanded octets

As mentioned previously, the octet rule works best for the elements in the second period (Li through F) of the periodic table. The reason for this is that electrons, whether shared or not, must be contained in orbitals, and the energies of electrons in such orbitals must be relatively low; otherwise, there would be no energetic advantage in forming a bond in the first place— the atoms would be better off by themselves.

Figure 7.58 Lewis Dot structures

For the second- and third-period elements, the n=2 and n=3 s²p⁶ sets comprise the octet. Some of the third-period elements (Si, P, S, and Cl) can bond to more than four atoms, and thus need to involve more than the four pairs of electrons available in an s²p⁶ octet. This is possible because at n=3, an additional set of d orbitals can exist. Although their energies are higher (ordinarily higher than the 4s orbitals), they can participate in the valence shells of these atoms.

Examples of molecules in which the n=3 central atom contains an expanded octet are the phosphorus pentahalides and sulfur hexafluoride.

Similarly, for atoms in the fourth period and beyond, higher d orbitals can sometimes be used to accommodate additional shared pairs beyond the octet . If you review a diagram showing the relative energies of the different kinds of atomic orbitals (see here, for example), you will notice that all the energy gaps become smaller as the principal quantum number increases, so the energetic cost of using these higher orbitals becomes smaller.

A multivalent molecule is a molecule that has atoms that can violate the octet rule and have more than eight electrons in their valence shells when these atoms are combining to form a covalent bonding.

Examples of these multi valent also called hypervalent molecules are given below:

Phosphorus pentachloride (PCl₅)

sulfur hexafluoride (SF₆)

chlorine trifluoride (ClF₃)

triiodide ion (I₃⁻)

The elements in the second period of the periodic table (with the principal energy level n = 2) have the valence electrons orbitals s²p⁶ and there are no d orbitals available and hence these elements strictly obey the octet rule without any deviations. These elements are nonmetals i.e. C, N, O and F.

With the principal energy level n = 3 elements such as Si, P, S and C can bond more that four atoms to make a covalent bond and therefore such elements can violate the octet rule.

The reason behind this violation or exception to the octet rule is that these elements have d sublevels. These d sublevel or sub orbitals are empty and have higher energy that the 4s orbitals in general but in the reality the difference between 3d and 4s energies are small and hence the 3d orbitals can accommodate more electrons.

The 3d orbitals do participate in the forming the covalent bonding and in this process the octet rule is expanded or violated.

Below are examples of such expansion or violation:

Reference: https://courses.lumenlearning.com/boundless-chemistry/chapter/exceptions-to-the-octet-rule/

The pm is a unit of measuring the length of the bond in Pico meter (10 ^-12).

Both phosphorous and sulfur are violating the octet rule:

Multiple bonds

If one pair of electrons shared between two atoms constitutes a chemical bond, it seems logical that two or three pairs could be shared to produce double and triple bonds. Such formulations appear quite naturally when the octet rule is applied to elements such as C, O, S, and N.

Figure 7.59 Multiple bonds formation

Ref: commons.wikimedia.org

The colors of the electron-dots in this illustration are intended only to help you keep track of the number of electrons contributed by each atom. Because electrons are indistinguishable, it makes no sense to identify a given electron-dot with a given atom.

Since multiple bonds place more electron density between the two nuclei, the latter are held toward each other more closely and tightly; multiple bonds are therefore shorter and stronger than single bonds.

Can there be bonds of higher order than three? No one thought so for a long time, but beginning in the 1960s, experiments and theoretical computations began to reveal that this just might be possible; some molecules just don’t follow the rules!

Adding to the fun, a 2006 article presented evidence that the carbon atom in C(PPh₃)₂ (where Ph stands for a benzene ring) has two lone pair electrons, but no electrons connected to the bonded groups— thus introducing the concept of a “zero-uple” bond.

PCl₅: Phosphorous carries 10 valence electrons combined                    SF₆: Sulfur carries 12 valence electrons combined

Figure 7.60 Exception to octet rule

Ref: commons.wikimedia.org