It shows unique physical and chemical properties due to its molecular structure and bonding. Molecular formaldehyde is a colorless, toxic gas with an irritating and pungent smell. In its solid form, it can exist either as a trimer (1,3,5-trioxane) or a polymeric form is known as Paraformaldehyde. Formaldehyde is usually stored in aqueous solutions known as formalin. This chemical, due to its high reactivity is also one of the most important building blocks of synthetic chemistry. One of the biggest clues to this versatility in the physical state as well as the reactivity of formaldehyde lies in the nature and structure of its chemical bonds. A good starting point in understanding them will be the Lewis structure.
H2CO Lewis Structure
A Lewis structure representation is the simplest way to indicate the chemical formula of a compound by showing its valence and bonding electrons along with the formal charge. The lone pair electrons are shown as dots on the atom they belong to whereas the bonding electrons are either shown as lines connecting the two atoms sharing them or as pairs of dots between them. The formal charge must always be mentioned in a Lewis structure. In the most complete structures, the formal charge around every atom is mentioned, in simpler ones only the overall charge on the molecule is written as a superscript outside square brackets.
How to Draw the Lewis structure of Formaldehyde
Step 1: Calculate total no. of valence atoms in molecule i.e. the group no. of every atom plus total negative charge or minus the total positive charge. Formaldehyde is a neutral molecule so it has zero net charges. Step 2: Choose the central atom. It is usually the atom able to form maximum no. of single bonds and/or able to expand its octet. In this case, it is carbon, which can form 4 bonds. Now join other terminal atoms using single bonds to the central atom. Step 3: Every single bond has used up 2 electrons, so the rest are added as lone pairs to the terminal atoms to complete the octets, till there are no more electrons left. It is a good idea to start with the most negative elements first while adding lone pair of electrons. We have formed three bonds using 6 electrons for formaldehyde. That leaves 6 electrons, all of which are used upon oxygen.
Step 4: We have to complete the octet on the central atom. If there are no electrons left, a lone pair from a terminal atom can be used to form another bond. In this case, from oxygen, we draw in a lone pair to complete the octet around carbon. Step 5: Now we calculate the formal charge on each atom, Formal charge=No. of valence electrons- 12No. of bonding electrons-No. of lone pairs Carbon: Formal charge = 4 –½ 4 – 0 = 0 Oxygen: Formal charge = 6 – ½4 – 4 = 0 Hydrogen: Formal charge = 1 – ½2 – 0 = 0 So, the final Lewis structure, with zero formal charges is:
H2CO Hybridization
We have the basic picture of bonding in the Lewis structure of H2CO but we still do not know about the shape of the molecule. Covalent bonds are directional which means they have a specific arrangement in space. Hybridization helps us understand the nature of these covalent bonds using atomic orbitals of the central atom. In the case of formaldehyde, let’s take carbon. Carbon has an electronic ground state configuration of 1s2 2s2 2p2. Unhybridized carbon will only be able to make 2 single bonds along the internuclear axis (usually the z-axis), using s orbital and one pz orbital. But there are three single bonds in Formaldehyde so hybridization becomes important.
The sp2 hybrid orbitals are planar with an angle of 120⁰. The unhybridized 2p orbital on carbon is perpendicular to the molecular plane. It is used for sideways overlap to form a π bond with the 2p orbital of oxygen. Oxygen is also sp2 hybridized in this molecule, but it forms only one σ bond as the other two sp2 hybrid orbitals are filled with lone pair electrons. We will see why the bond angles are not exactly 120⁰ in the next section. Hybridization can also be calculated for a molecule using the formula X = ½ * ( H + V + A – C) where V = No. of valence electrons in the central atom H = No. of monovalent terminal atoms C = cationic charge A = anionic charge X = No. of hybrid orbitals For formaldehyde, X = ½ * ( 2 + 4 + 0 + 0 ) = 3. Three hybrid orbitals can be formed from one s and two p orbitals giving sp2 hybrids.
H2CO Molecular Geometry
The Valence Shell Electron Repulsion Theory attempts to predict the geometry of individual molecules using the concept of minimum energy and maximum stability. According to VSEPR, the lowest energy can be achieved by minimizing repulsion between electron pairs around the central atom, giving the most stable geometry. In formaldehyde, we will be considering the electron pairs around Carbon. To apply VSEPR, we need the steric no. of Carbon which is the no. of atoms bonded to the central atom plus the no. of lone pair electrons on the central atom. For carbon, it is three. According to the table below, for total domains (steric no.) = 3 and lone pair = 0, the molecular shape is trigonal planar.
According to VSEPR, the magnitude of repulsion increases in the order: Bond pair- bond pair < Lone pair- bond pair < Lone pair- Lone pair The double bond between carbon and oxygen also offers more repulsion than the carbon-hydrogen single bond, distorting the molecule slightly. The O-C-H bond angles up to > 120⁰ and the H-C-H bond angle closes to < 120⁰.
H2CO Molecular Orbital (MO) Diagram
So far, we have covered many structural aspects of the formaldehyde molecule by using hybrid orbitals and VSEPR theory. In addition, we have a molecular orbital theory (MOT), the most in-depth analysis of chemical bonding using quantum mechanical properties. MOT describes electron density and distribution in molecules similar to atomic orbitals in atoms. It uses a wavefunction Ψ to describe the behavior of valence electrons, molecular orbitals are filled up in the same way as atomic orbitals and the total no. of molecular orbitals formed = no. of combining atomic orbitals. The molecular orbital diagram represents energy on the Y axis, with bonding molecular orbitals lower in energy than anti-bonding orbital. A bonding molecular orbital represents electron density in between the two nuclei where attraction overcomes repulsion. An anti-bonding orbital represents the region between two nuclei where repulsion is greater than attraction and no bond is formed. Non-bonding orbital represents atomic orbitals which essentially remain localized on a single atom and do not take part in bond formation. When there are more than two atoms involved in bond formation, an approximation method called a linear combination of atomic orbitals (LCAO) is used, like in formaldehyde. MOT helps us determine electrostatic potential maps, excited state properties, probable molecular spectra, and predict the reactivity of the molecule.
Conclusion
In this article, we have so far learned about the steps to draw the lewis structure of the H2CO molecule and about the formation of molecular geometry using VSEPR theory with its molecular orbital theory as well. If you have any questions regarding this topic or any suggestions, feel free to comment below. I’ll be glad to read your comments.
