How to Draw Lewis Structures

The electron density around atoms is depicted graphically in a Lewis structure. Learning to sketch Lewis structures allows you to anticipate how many and what kind of bonds will form around an atom. A Lewis structure can also be used to predict the geometry of a molecule. The models often confuse students, but if the appropriate processes are followed, sketching the Lewis structures can be a simple task. Be aware that there are a variety of approaches to building Lewis dot structures. The Kelter approach for drawing Lewis structures for molecules is outlined in these instructions.

Step 1. Finding the Total Number of Valence Electrons is the first step.

Add the total number of valence electrons from all the atoms in the molecule in this stage.

Step 2: Determine how many electrons are required to make the atoms “happy.”

When an atom’s outer electron shell is filled, it is said to be “happy.” Eight electrons are required to fill the outer electron shell of elements in the periodic table up to period four. The “octet rule” is a term used to describe this characteristic.

Step 3: Work out how many bonds there are in the molecule.

When one electron from each atom forms an electron pair, covalent bonds are established. Step 2 informs you how many electrons you’ll need, and Step 1 tells you how many you already have. The number of electrons required to complete the octets is calculated by subtracting the number in Step 1 from the number in Step 2. Because each bond requires two electrons to form, the total number of bonds is half of the total number of electrons required.

Step 4: Select a Centre Atom

The least electronegative atom in a molecule, or the atom with the highest valence, is usually the central atom. Either depend on periodic table trends or use a table that gives electronegativity values to find electronegativity. Travelling down a group on the periodic table decreases electronegativity, while moving from left to right across a period increases it. Hydrogen and halogen atoms are usually found on the molecule’s periphery and are rarely found in the centre.

Step 5. Draw a Skeletal Structure in Step 5

A straight line indicating a link between the two atoms connects the atoms to the central atom. Up to four additional atoms can be joined to the core atom.

Step 6. Place Electrons Around Outside Atoms in Step 6

Fill in the octets all the way around each of the outside atoms. The skeletal structure from Step 5 is wrong if there aren’t enough electrons to complete the octets. Attempt a new arrangement. This may necessitate some trial and error at first. It will become easier to forecast bone structures as you gain experience.

Step 7. Place the remaining electrons around the central atom in step 7.

With the remaining electrons, complete the octet for the centre atom. Create double bonds with lone pairs on exterior atoms if any bonds remain from Step 3. Two solid lines are painted between two atoms to indicate a double bond. 

If the centre atom has more than eight electrons and is not one of the octet rule’s exceptions, the number of valence atoms in Step 1 may have been wrongly counted. This will complete the molecule’s Lewis dot structure.

Bonding atoms in Lewis structure : 

Molecular bonding and lone electron pairs are depicted in Lewis structures, also known as Lewis dot formulas, Lewis dot structures, electron dot structures, or Lewis electron dot structures (LEDS) diagrams. A Lewis structure can be used to represent any covalently attached molecule, as well as coordination compounds. The Lewis structure was first described in Gilbert N. Lewis’ 1916 article The Atom and the Molecule, and it was named in his honour after him. Adding lines between atoms to represent shared pairs in a chemical bond extends the concept of the electron dot diagram, which is an extension of the concept of Lewis structures.

Number of bonds in Lewis structure : 

It is equal to the number of electrons in the complete valence shell (2 or 8 electrons) of a neutral atom, minus the number of valence electrons, to determine the number of bonds for a neutral atom. This strategy works because each covalent link that an atom makes increases the number of electrons in an atom’s valence shell without affecting the charge of the atom involved in the bond.

Pairs in bonding atoms : 

Each chlorine atom has three pairs of electrons, two of which are shared by the other atoms. These are referred to as lone pairs, and the other two electrons on each chlorine atom are known as bonding pairs. Lone pairs are not involved in the formation of covalent bonds. The bond is referred to as a coordinate covalent bond if both electrons in a covalent bond originate from the same atomic nucleus.

Resonance in Lewis structure : 

It can be difficult to decide which lone pairs should be relocated in order to make double or triple bonds in some molecules and ions, and it is possible to write two or more distinct resonance structures for the same molecule or ion in some cases. If there are several of them, it is customary to write them all with two-way arrows in between. Occasionally, this is the case when a core atom is surrounded by numerous atoms of the same type; this is especially typical in the case of polyatomic ions.

When this occurs, the Lewis structure of the molecule is said to be a resonance structure, and the molecule is said to exist as a resonance hybrid. It is assumed that the molecule has a Lewis structure that is comparable to some combination of the numerous possibilities since each of the many possibilities is overlaid on the others.

For example, the nitrate ion (NO3-) must establish a double bond between nitrogen and one of the oxygens in order to meet the octet rule for nitrogen in its structure. However, because the molecule is symmetrical, it makes no difference which of the oxygens is responsible for forming the double bond in this case. There are three different resonance structures that might be used in this situation. When sketching Lewis structures, it is possible to express resonance in two ways: by drawing each of the possible resonance forms and connecting them with double-headed arrows, or by using dashed lines to indicate 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, those with the fewest formal charges tend to contribute more to the overall resonance hybrid than those with the most formal charges do. When formal charges are required, resonance structures with negative charges on the more electronegative components and positive charges on the less electronegative elements are preferred over other types of resonance structures.

As a result, single bonds can be manipulated in the same way to form resonance structures for hypervalent compounds such as sulphur hexafluoride, which is the right description according to quantum chemical calculations rather than the often used expanded octet model.

The resonance structure should not be understood as indicating that the molecule flips between different forms, but rather as indicating that the molecule works as an average of numerous different forms.

Conclusion : 

The electron density around atoms is depicted graphically in a Lewis structure. Learning to sketch Lewis structures allows you to anticipate how many and what kind of bonds will form around an atom.It can be difficult to decide which lone pairs should be relocated in order to make double or triple bonds in some molecules and ions, and it is possible to write two or more distinct resonance structures for the same molecule or ion in some cases.