The Geometry of the Orbital Arrangement

Molecular geometry refers to a molecule’s three-dimensional arrangement of atoms. A molecule’s molecular geometry, or form, is a significant aspect that influences a compound’s physical and chemical properties. Melting and boiling points, solubility, density, and the types of chemical reactions that a compound undergoes are all examples of these qualities. In this course, you’ll learn how to use the Lewis electron dot structure of a molecule to predict molecular geometry.

VSEPR Theory

The valence shell is an atom’s outermost occupied electron shell. The valence electrons, which are involved in bonding and illustrated in a Lewis structure, are held in this shell. According to the valence-shell electron pair repulsion model (VSEPR), a molecule would change its structure to keep the valence electron pairs as far away as feasible. Based on the fact that negatively charged electrons repel one another, this makes sense. The number of bonding pairs of electrons and the number of nonbonding or lone pairs around the core atom will be used to define molecules. In the VSEPR model, a double or triple bond is no different from a single bond in terms of repulsion. We’ll start by looking at compounds in which the core atom has no lone pairs.

Linear

A bond angle is essentially the geometric angle between two neighbouring bonds in a linear model, where atoms are joined in a straight line. The two bonding orbitals of a simple triatomic molecule of type AX2 are 180 degrees apart. BeCl2 (which lacks enough electrons to comply with the octet rule) and CO2 are examples of triatomic molecules for which VSEPR theory predicts a linear shape. Notice that the C-O bonds are double bonds when putting out the electron dot formula for carbon dioxide; this makes no difference to VSEPR theory. Two other atoms are still connected to the centre carbon atom. The electron clouds that connect the two oxygen atoms are at a 180-degree angle.

Trigonal planar

The trigonal planar shape refers to molecules that are triangular and have only one plane, or flat surface. Three zones of electron density spread out from the centre atom in an AX3 molecule like BF3 . When the angle between any two is 120 degrees, the repulsion between them is at a minimum.

Tetrahedral

“Tetrahedral” literally means “having four faces,” as tetra- denotes four and -hedral denotes a solid face. This structure is generated when a single core atom has four bonds and no lone electron pairs. The bond angles between the electron bonds are 109.5o, according to the VSEPR theory. Methane is an example of a tetrahedral molecule (CH4).The four analogous bonds, which correspond to the four corners of a tetrahedron centred on the carbon atom, point in four geometrically equivalent directions in three dimensions.

Trigonal bipyramidal

In chemistry, a trigonal bipyramid formation has one atom at the centre and five other atoms in the corners of a triangular bipyramid. Because there is no geometrical configuration with five terminal atoms in equivalent places, the bond angles surrounding the central atom are not identical (see also pentagonal bipyramid). In the gas phase, phosphorus pentafluoride (PF5),and phosphorus pentachloride (PCl5) are examples of this molecular geometry.  When a core atom is surrounded by five atoms in a molecule, it forms a trigonal bipyramidal structure.Three atoms are on the same plane with bond angles of 120° in the geometry, while the other two are on opposite ends of the molecule. PCl5 and AsF5 are two examples of AX5; compounds formed by elements in Group 15 of the periodic table.

Octahedral

In chemistry, octahedral molecular geometry describes the form of compounds with six atoms or groups of atoms or ligands symmetrically arranged around a central atom, defining the vertices of an octahedron. The octahedron is named after its eight faces, hence the prefix octa.Although octahedral compounds typically have a single atom in the centre and no links between the ligand atoms, the octahedron is one of the Platonic solids. Sulfur hexafluoride SF6 and molybdenum hexacarbonyl Mo(CO)6 are examples of octahedral chemicals. Chemists use the word “octahedral” loosely, concentrating on the geometry of the bonds to the central atom rather than distinctions among the ligands themselves. [Co(NH3)6]3+,is a good example. 3+ is referred to as octahedral, despite the fact that it is not octahedral in the mathematical sense due to the orientation of the NH3 bonds. Because octa- denotes eight and -hedral denotes a solid face, “octahedral” literally means “having eight faces.” The bond angles are all 90 degrees, and six electron pairs try to point toward the corners of an octahedron, just as four electron pairs feel least repulsion when directed toward the corners of a tetrahedron. Sulfur hexafluoride is an example of an octahedral molecule(AX6) (SF6).

Conclusion:-

Hybridization of orbitals is preferred because hybridised orbitals are more directed, resulting in more overlap when forming bonds, resulting in stronger bonds. When hybridization occurs, this results in more stable molecules.