The equation for the preparation of Oxygen Difluoride: 2F2 + 2NaOH ——> OF2 + 2NaF + H2O It is a colorless gaseous compound exhibiting a strong peculiar odor and acts as an oxidizer. It has a boiling point of 128.40 K and a melting point of 49.3 K. It has a molar mass of 53.996 g/mol and a density of 1.88 g/l as a gas at room temperature. Although it has its use as a strong oxidizing agent, for example, in propellants of rocket fuels, this property also makes the compound unsafe. OF2 can explode on contact with water and is said to be highly corrosive to the skin. It can also affect the eyes and the cardiovascular system. Let us have a quick but detailed look into the inside of the molecule and discuss its nature of chemical bonding.
OF2 Lewis Structure
Do you know that several factors make atoms come together and combine to form several different chemical compounds? When atoms come together and create bonds, a new molecule is created. Usually, the electrons in the outermost shell of an atom, also known as valence electrons, take part in chemical bonding. Lewis Structure, thus, is a simple and constructive procedure to sketch a diagrammatic representation of a molecule with the help of electron-dot structures. It gives us a 2D figure and helps us have a brief idea about the electron arrangement and the type of bond formation without going much into detail. Now, we will discuss the steps to form the Lewis Structure of OF2.
Steps to form OF2 Lewis Structure Diagram
Step 1: Find the Total number of Valence Electrons. The first and foremost step is to calculate the total number of valence electrons in an OF2 molecule. Oxygen belongs to group 16, the chalcogen family, and has a valency of 6. Fluorine belongs to the family of halogen in group 17 and has a valency of 7. Therefore, the total number of valence electrons = 6 + 7*2 = 20. Step 2: Identify the Central Atom Now, we have to decipher the central atom in this molecule. If we look at the Pauling Electronegativity chart, we can find out that the least electronegative element among oxygen and fluorine is Oxygen (O).
The central atom in an oxygen difluoride molecule is Oxygen. Step 3: Sketch the Skeletal Diagram of the Molecule In Lewis Structure, we use atomic symbols like C for carbon, H for hydrogen to represent the constituent atoms, and electron dot notation to represent the valence shell electrons. Let us look at the below skeletal sketch:
The atomic symbols:
Atomic symbols along with dot notations:
Step 4: Octet rule The main group elements ( group 1 to group 17 ) of the periodic table have a tendency to attain the octet configuration of the noble gas elements present in group 18 of the same period. This is known as octet fulfillment since the elements want to achieve eight electrons in their valence shells. For example, Carbon tends to attain a Neon configuration. Exception: Hydrogen tends to achieve Helium configuration, hence only two electrons in the outer shell. Here, if we look at the octet rule, both the fluorine atoms have achieved a valency of eight. The same goes for the central oxygen atom. Step 5: Formal Charge Concept Before we proceed to confirm any sketch to be the perfect Lewis Structure for a given molecule or ionic structure, we need to check the formal charge values.
For Oxygen: Formal Charge = 6 – 0.54 – 4 = 6 – 2 – 4 = 0. For each Fluorine atom: Formal Charge = 7 – 0.52 – 6 = 7 – 1 -6 = 0. Since the elements are present in their least possible formal charge values, we have achieved our suitable Lewis Structure configuration. Let us see:
Here, oxygen forms a single bond with each of the fluorine atoms (O-F).
OF2 Molecular Geometry
We have already found the 2D Lewis Structure diagram of the Oxygen Difluoride molecule. Now, we are going to decipher the 3D molecular shape. Via Lewis Structure, we have realized the type of bond formed and the number of lone or unbonded pairs of valence electrons present in an OF2 molecule. Our next step is to use Valence Shell Electron Pair Repulsion Theory (VSEPR) to find the exact molecular geometry along with the bond angles. Like charged electrons create a negatively charged cloud surrounding the atomic nuclei and repulsion is created. The VSEPR model tends to reduce or minimize the repulsive forces and balance the stability of any given molecule, in this case, OF2. Let us look at this VSEPR chart showing AXN notations:
According to AXnEx notation of VSEPR theory, A: central atom ( Oxygen ) X: no. of surrounding atoms, ‘n’ = 2 ( Fluorine ) E: no. of lone pairs of electrons of the central atom, ‘X = 2. The VSEPR notation for the OF2 molecule is AX2E2. According to the above chart, we have a bent molecule structure. The bond angle is around 103 degrees (due to the repulsion of the lone pair).
Hybridization
In chemistry, an orbit is a fixed path around the nucleus in which electrons tend to move or revolve. An orbital, on the other hand, is defined to be the probability of the presence of an electron in any given regional space. Atomic orbitals are of several shapes like spherical, dumb-bell, clover-leaf, or doughnut to name a few of the complex varieties. We have s,p,d,f atomic orbitals ( AOs ). Electrons can be explained via their wave functions in quantum mechanics. If we take into account the combination of wave functions of several atomic orbitals, we come across the process of hybridization. For example, the wave functions of the s and a 2px orbitals give rise to sp hybridization which usually results in a linear geometry.
Hybridization in OF2
Let us look at the electronic configuration of oxygen and fluorine in their ground states. O: 1s2 2s2 2p4 F: 1s2 2s2 2p5
After forming bonds with halogen F and attaining octet configuration, the electronic structure of O looks like this:
We, therefore, have in the outermost shell of the central oxygen atom, four hybridized orbitals: 2s, 2px, 2py, 2pz. Steric number = Number of atoms bonded to central atom inside a molecule + Number of lone pair of electrons attached to the central atom Also, steric number = 2 + 2 = 4. Therefore, we have sp3 hybridization for the central oxygen atom in a molecule of oxygen difluoride ( OF2 ).
OF2 Polarity
Let us understand at first what we mean by the very famous and common terminology: Polarity. Polarity can be defined as the physical property of any molecule or chemical compound where the constituent atomic elements can possess either positive or negative charges or both. We have already seen the Pauling Electronegativity chart. Each and every element of the main groups in the periodic table has a certain value of electronegativity. When atoms of the same element come together in equal ratio, the charge values neutralize themselves and the net dipole is zero, therefore the molecule is non-polar. Linear or symmetrical chemical molecules are usually non-polar. ( e.g. : H2 ) However, if molecules consist of uneven partial charge distributions due to a difference in electronegativity or the molecular shape is asymmetrical, the net dipole doesn’t cancel or becomes zero and we have a polar molecule. ( e.g. : H3PO4, HCl ). You can also refer to the polarity of HCl.
Is OF2 a polar molecule?
Let us have a quick look into the molecular shape of Oxygen Difluoride. It is almost similar to H2O and in both cases, the central O has two lone pairs on the top. The electronegativity value of Oxygen is 3.44. And, that of fluorine is 3.98. Since oxygen is more electropositive than fluorine atoms, it will possess a partial positive charge ( δ+ ) and each of the F atoms will possess a negative partial charge ( δ- ). So, we can see that there is an electronegativity difference and henceforth the net dipole moment is not equal to zero. Again, OF2 is not a symmetrical linear structure. The lone pairs on the central Oxygen atom give it a bent molecule geometry which resembles the English letter V. We have gone through both the uneven charge distribution and the asymmetry of Oxygen Difluoride. The molecule is polar in nature. For detailed information, you must also read out the article written about the polarity of OF2.
OF2 Molecular Orbital (MO) Diagram
Valence Bond Theory is important for explaining chemical bonding inside a molecular structure but it has several limitations or drawbacks. Here, comes the Molecular Orbital Theory or MOT. In MOT, we take into account the concept of molecular orbitals (MOs) which are orbitals of the molecule as a whole and not just those of constituent individual atoms. It considers electrons to be delocalized and results in creating bonding and anti-bonding interactions in orbitals. To showcase the relative energy levels of the AOs and the resultant MOs, we have the Molecular Orbital diagrams. The lower energy molecular orbital is the bonding and the higher energy is the anti-bonding orbital. The below-mentioned diagram gives us the individual MO diagrams of Oxygen and Fluorine separately, the atoms that make up a molecule of OF2.
Conclusion
In this article, we have discussed the type of chemical bonding occurring inside an OF2 molecule. Oxygen Difluoride is an interesting molecule that was discovered accidentally and later had several applications although it was toxic in nature. We have covered Lewis Structure, VSEPR theory to determine molecular geometry, Hybridization, Polar nature, and a brief discussion on MO diagram about oxygen difluoride. Happy learning!
