The first step in visualizing a molecule is to convert its molecular formula to its Lewis structure (or Lewis formula*), which depicts atom symbols, bonding electron pairs as lines, and lone electron pairs that occupy each atom's outer level (valence shell) as pairs of dots.

To write a Lewis structure, we first decide on the relative position of the atoms in the molecule or polyatomic ion, and then we distribute the total amount of valence electrons as bonding and lone pairs.

The octet rule (Section 9.1) assists us in dispersing electrons in many, but not all, circumstances. Remember that shared (bonding) electrons are tallied as part of each atom's octet, but lone pairs are counted exclusively by the atom to which they belong.

We start with species that "follow" the octet rule, in which each atom has eight electrons on its outer level (or two for hydrogen).

You may write a Lewis structure for every singly bonded species with a central C, N, or O atom, as well as certain species with central atoms from higher periods, using these four stages: Hydrogen atoms create one link in virtually all of their compounds.

Carbon atoms make four bonds, nitrogen atoms have three connections, and oxygen atoms are two bonds. Surrounding halogens create a single bond; fluorine is always a halogen atom.

A stepwise process converts a molecular formula into a Lewis structure, a two-dimensional representation of a molecule (or ion) that shows the placement of atoms and the distribution of valence electrons among bonding and lone pairs.

The actual structure is a hybrid of those resonance structures when two or more Lewis structures can be drawn for the same relative placement of atoms.

Formal charges can be beneficial in determining which contributor to the hybrid is more important, but experimental data always determines the option.

Molecules containing an electron-deficient atom (central Be or B) or an odd-electron atom (free radicals) have less than an octet but frequently achieve an octet.

Almost every biological activity is heavily dependent on the morphologies of interacting molecules.

Every medication you take, every odor you smell, and every flavor you taste is dependent on a portion of one molecule fitting together with another.

Complex behaviors in many species, such as mating, defense, navigation, and eating, have been discovered by biologists to rely on one molecule's structure matching that of another. In this part, we will look at a model for forecasting a molecule's shape.

Chemists begin with the Lewis structure and then apply the following theories to transform that two-dimensional structure into a three-dimensional shape to establish the molecule shape:

Valence-shell electron-pair repulsion (VSEPR) theory states that in order to reduce repulsions, each group of valence electrons around a core atom should be as far apart as possible from the others.

A "group" of electrons is any number of electrons that occupy a confined area surrounding an atom.

A single electron group might consist of a single bond, a double bond, a triple bond, a lone pair of electrons, or a single electron.

Only electron groups around the core atom have an effect on form; electrons on atoms other than the central atom have no effect.

The three-dimensional arrangement of nuclei linked by bonding groups is known as the molecular shape.

When two, three, four, five, or six items connected to a central point maximize the space between them, five geometric patterns emerge, as seen in Figure 10.2A with balloons.

If the objects are valence-electron groups, repulsions maximize the area each occupies around the center atom, resulting in the five electron-group configurations found in the vast majority of molecules and polyatomic ions.

The bond angle is the angle created by the bonds connecting the nuclei of two surrounding atoms to the center atom's nucleus, which is at the vertex.

The angels depicted in the image attached are ideal bond angles defined only by fundamental geometry.

When all of the bonding groups are the same and linked to the same sort of atom (structure I, below), we can see them. In the absence of this, the actual bond angles vary from the ideal angles.

Deviations arise when the bonds are different (structure II), the surrounding atoms are different (structure III), or one or more of the electron groups are nonbonding groups (structure IV).