Shapes of Some Simple Molecules

Introduction:

The geometries of molecules, as well as their bond angles (the angles between covalent bonds), are controlled by a set of concepts known as Valence Shell Electron Pair Repulsion Theory. (VSEPR)

A simple molecule is made up of a centre atom and a number of additional atoms organised around it. This centre atom’s outer shell will contain electrons that are part of bonded pairs, as well as lone pairs. The molecule’s structure and bond angles are determined by the mutual repulsion between these electron pairs.

According to VSEPR, all electron pairs repel each other and will endeavour to take up positions as far apart as feasible to minimise repulsion. It merely relies on our understanding of electron behaviour to predict the structure of basic covalent molecules. To refresh your memory on how atoms share electrons to form stable electron configurations, go to Covalent and Dative Bonding.

Drawing shapes of molecules in 3D:

Before we look at any examples of covalent structures, we need to understand how to represent them in three dimensions. Covalent bonds can be drawn as a line between two atoms, as you may recall. This depicts molecules in a straightforward manner. We can utilise wedge and dotted lines, though, to better represent the 3D form of a molecule.

Wedged lines depict a relationship emerging from the screen or page and moving towards you.A bond travelling into the screen or page away from you is shown by dotted or dashed lines.Dots represent lone pairs of electrons.A planar link can be seen in any ordinary straight line.

Molecules come in a variety of shapes:

If all of an atom’s valence electron pairs are linked together, they will reject each other. As a result, the bonds are evenly spaced. The shape of the molecule and the angle between the bonding pairs are affected by the number of bonded electron pairs.

Let’s discuss the most popular shapes. Keep in mind, however, that these laws only apply to compounds that lack lone pairs of electrons. Unshared pairs of electrons that aren’t covalently bound are known as lone pairs. Later, we’ll go deeper into their impact.

Linear:

A linear molecule is formed when a molecule has just two bound electron pairs (and no lone pairs). Beryllium chloride is the most basic example. Despite being a metal, beryllium may form a covalent connection with chlorine. Beryllium’s valence shell only has two electrons, hence it forms two bonds. The electron pairs reject each other in an equal amount, resulting in an angle of 180° between the two bonds.

Planar trigonometry:

Trigonal planar molecules are those that have three bound electron pairs. Because each bond has a bond angle of 120 degrees, the bonds lie level on a plane. The molecules could be stacked one on top of the other like sheets of paper. One example is boron trifluoride.

Tetrahedral:

Tetrahedral molecules are those that have four bound electron pairs but no lone pairs. This pyramid has a standard triangular base. The bond angles are all 109.5 degrees. The carbon atom in methane (CH4) has four valence electrons, each of which is part of a pair that is covalently bound to a hydrogen atom. The molecule is tetrahedral.

Trigonal bipyramid:

A trigonal bipyramid is made up of molecules having five bound electron pairs. This molecule has a trigonal planar form with two additional bonds extending above and below the plane at 90 degrees. A good example is phosphorus(PCl5) pentachloride.

Octahedral :

An octahedral structure is formed when a molecule has six bonding pairs around a central atom. As demonstrated in sulphur hexafluoride(SF6), all of the bonds are at right angles to one another.

Lone pairs of electrons: 

All of the examples above involve molecules with no lone pairs of electrons. Their valence electrons are all bonded together.But what if a molecule has only one lone pair. As an example, consider a molecule with four electron pairs.

We now know that the molecule will be tetrahedral and have bond angles of 109.5° if all of the electrons are part of bonding pairs. The bond angles are reduced to 107° if one of the electron pairs is actually a lone pair. This is due to the fact that lone pairs repel each other more strongly than shared pairs, compressing the bonds.

Pyramidal: 

The angle between each bond in a molecule with three bonded electron pairs and one lone electron pair around a central atom is 107°. Ammonia is a good example. In nitrogen there are five valence electrons. The remaining two form a lone pair and are covalently bonded to hydrogen atoms.

V-shaped:

The bond angle of a molecule with two lone pairs and two bonding pairs is lowered to 104.5° when it is V-shaped. This results in the formation of a v-shaped molecule, such as water.

Conclusion: 

VSEPR, or valence shell electron pair repulsion theory, states that electron pairs balance each other and aim to take up locations as far apart as possible to minimise repulsion. This has an impact on the structures of molecules. Straight lines are used to symbolise covalent bonding. Wedged lines indicate a bond that extends from the page, whereas dashed or dotted lines indicate a bond that extends backwards.