Dichloromethane or methylene chloride, with the chemical formula CH2Cl2, is a colorless, volatile liquid with a boiling point of 39.6 °C. and a melting point of -96.7 °C. It is widely used as a solvent in chemistry laboratories.
It is polar because of the presence of two chloro groups but is not miscible with water; however, it does show miscibility with various organic solvents such as chloroform, carbon tetrachloride, hexane, benzene, ethyl acetate, and alcohols.
The preparation of CH2Cl2 involves a high-temperature treatment of methane or chloromethane with chlorine gas.
CH2Cl2 is considered toxic; its overexposure via inhalation leads to dizziness, nausea, numbness, and weakness. It is also metabolized in the body to form carbon monoxide and can cause poisoning.
CH2Cl2 Lewis Structure
The Lewis theory of chemical bonding—although quite primitive and the most limited theory on electronic structure—does help one to determine how valence electrons are arranged around the constituent atoms in a molecule.
The purpose of this theory is to help visualize the chemical bonding of atoms in molecules.
Electrons are represented as dots, and each pair of bonding electrons between two atoms is shown as a line. The structures drawn using this theory are termed Lewis (dot) structures.
Please note that several atoms follow the octet rule, i.e., they tend to achieve eight electrons in their valence shell through chemical bonding; this is reflected in the Lewis structure of the molecule.
Hydrogen, however, does tend towards a duplet, not octet, because it has only one electron in its K shell, and thus needs only one more to achieve the maximum capacity of the K shell.
Let us take a look at the chemical bonding represented by Lewis structure in CH2Cl2.
Step 1. We shall start by calculating the number of valence electrons in each atom of CH2Cl2 in order to see how short an atom is from an octet (or duplet in the case of hydrogen).
i. The atomic number of carbon is 6; therefore, it possesses 6 electrons in its neutral form. There are 2 electrons in its K shell and 4 electrons in the L shell. Thus, the number of valence electrons is 4. To achieve the octet, carbon needs 4 more electrons.
ii. Similarly, the atomic number of hydrogen is 1; thus, each H has 1 electron and needs 1 more to achieve the duplet.
iii. The atomic number of chlorine is 17. K shell has 2 electrons, L shell has 8, and M shell has 7 electrons. The number of valence electrons is therefore 7, and hence Cl needs 1 more to achieve the octet.
Step 2. Next, we shall figure out the central atom to which the rest of the atoms shall be bonded. The central atom is the one that has the highest bonding capacity; it is the atom that is the shortest of the octet. In CH2Cl2, carbon satisfies this condition (4 electrons short of the octet versus 1 for chlorine).
Step 3. Now, we shall construct a skeleton of the molecule with carbon as the central atom. Carbon needs 4 more electrons for its octet to be complete. Two hydrogen atoms and two chlorine atoms can help carbon achieve this feat!
Simultaneously, both hydrogen atoms will achieve their respective duplets, and both chlorine atoms will achieve their respective octets, and thereby the situation will be a win-win for all five atoms.
Carbon will be singly bonded to H, H, Cl, and Cl, as shown in the Lewis structure.
The molecule is neutral, i.e., there is no charge on it. Let us calculate the formal charges on each of the constituent atoms. The formula for the formal charge is as follows.
Formal charge (FC) = Valence electrons – 0.5*bonding electrons – non-bonding electrons
For carbon, FC = 0; for hydrogen, FC = 0; and for Cl, FC = 0.
A bond is formed between two atoms by the virtue of the overlap of orbitals on two atoms as these orbitals share electrons.
Let us look at the ground state electronic configuration of each atom in CH2Cl2 in terms of the orbitals.
Carbon, in the excited state, has one of the 2s electrons promoted to 2p; therefore, the electronic configuration becomes 1s22s22px12py12pz1. 2s, 2px, 2py, and 2pz orbitals of carbon are now half-filled.
These four orbitals hybridize together to form four identical sp3 orbitals, all of which have the same energy. Each of these hybrid orbitals has one electron and can accept one more.
One electron each comes from H, H, Cl, and Cl atoms: 1s1 of each H and 3pz1 of each Cl. This leads to the formation of four single bonds (also called sigma bonds) with four sp3 hybrid orbitals of carbon.
Another way of determining the hybridization of the central atom is by using the following formula.
Hybridization = A + (VE – V – C)/2,
A is the number of atoms/groups attached to the central atom;
VE is the number of valence electrons on the central atom;
V is the valency of the central atom;
C is the charge on the central atom.
|Value of H||
Type of hybridization
Herein, A = 4, VE = 4, V = 4, C = 0; therefore, Hyb = 4, corresponding to sp3.
CH2Cl2 Molecular Geometry
The geometry of a molecule can be determined using the hybridization of the central atom. Corresponding to sp3 hybridization, the geometry is tetrahedral when there are no lone pairs of electrons on the central atom.
Valence shell electron pair repulsion (VSEPR) theory helps to determine the geometry of a molecule on the basis of stoichiometry, the number of bond pairs, and the number of lone pairs on the central atom.
The fundamental idea behind this theory is that a molecule adopts such an arrangement of its constituent atoms that the repulsion arising from the valence shell electrons on all atoms is minimum.
The following table lists this information—on the basis of VSEPR theory— for various molecular stoichiometries.
According to the above table, the geometry of CH2Cl2 is tetrahedral, corresponding to the conditions stated for AX4.
The tetrahedral shape of CH2Cl2 is not perfect unlike that of CH4. This is because CH4 has all the identical hydrogen atoms around carbon, whereas CH2Cl2 has 2 H and 2 Cl.
This is reflected in the slight asymmetry in the molecular shape of the latter. This means that the bond angles and bond lengths in CH2Cl2 are not identical; however, all bond angles are identical in CH4.
The CH2Cl2 molecule is polar in nature.
As the shape of the molecule is tetrahedral and Carbon and Chlorine have a difference in their electronegativity. The asymmetric shape and electronegativity difference between atoms is an important aspect in determining whether a molecule is polar or not.
Thus C-Cl bond is polar and the overall charge distribution across the molecule is non-uniform. For detailed information, you must read out an article on the polarity of CH2Cl2.
Molecular Orbital Diagram for CH2Cl2
The premise of molecular orbital (MO) theory is that all the constituent atoms contribute towards the formation of molecular orbitals, which are a linear combination of the atomic orbitals. As per this theory, the electrons in a molecule are not individually assigned to atomic orbitals but to molecular orbitals.
Check out the MO diagram for CH2Cl2. The 2s and 2p orbitals of carbon mix (to different extents) with 1s orbitals of the two hydrogen atoms and 2pz orbitals of the two chlorine atoms.
Note that there are 8 atomic orbitals mixing to form 8 molecular orbitals. The extent of mixing and thus the contribution of individual atomic orbitals to form a particular molecular orbital depends on the relative energy alignment of the atomic orbitals.
Electron filling starts from the least energetic molecular orbital. The filled molecular orbitals are called bonding orbitals; the unfilled ones are anti-bonding orbitals.
When there is a lone pair an atomic orbital, that atomic orbital does not mix with any other orbital and forms a non-bonding molecular orbital.
The overview provided in this article helps in establishing a basic understanding of the structure of CH2Cl2 through chemical bonding.
In case, you have questions floating in your mind, please let me know.