H2O2 is a chemical compound with the IUPAC name Hydrogen Peroxide. It is the simplest peroxide compound, i.e., a molecule containing an Oxygen-Oxygen single bond.
It is a pale blue liquid in its standard state and slowly reacts with sunlight and decomposes into water and oxygen.
2H2O2 → 2H2O + O2 (in the presence of sunlight)
This is an example of a disproportionation reaction as O is a -1 state on the left-hand side while it exists in -2 and +0 oxidation states on the right-hand side.
H2O2 has a melting point of -0.43 Celsius (31.23 Fahrenheit). A low melting point indicates the tendency of the compound to remain in the liquid state.
It also exhibits a relatively high boiling point of 150.2 Celsius (302.4 Fahrenheit), attributed to strong hydrogen bonding interactions with water and other H2O2 molecules.
Industrially, H2O2 is prepared using Anthraquinone as a catalyst. The reaction is as follows :
2H2O + O2 → 2H2O2 (in the presence of Anthraquinone)
This is an example of a comproportionation reaction, i.e., the reverse of a disproportionation reaction.
It is a strong oxidizing agent, and hence, it finds wide applications as a bleaching agent and disinfectant.
It also plays a vital role in organic synthesis, used in the oxidation/reduction of various functional groups. It is also used as an oxidizer in spacecraft since oxygen is not available for combustion in outer space.
H2O2 is a toxic byproduct of several biological processes, and several enzymatic reactions are dedicated to the decomposition of H2O2.
It is toxic and corrosive and hence must be handled with caution.
H2O2 Lewis Structure
We shall discuss the chemical bonding nature of H2O2 in this article.
First, we shall draw the Lewis structure of H2O2.
A Lewis structure may not give us a complete description of the molecular geometry. Still, it is helpful to visualize the overall skeletal structure of the molecule, its bonds, and its lone pairs.
This is done by systematically analyzing the valence electrons of the constituent atoms.
Step 1: Calculate the total number of valence electrons in H2O2.
The periodic table shows that hydrogen (H) belongs to Group 1 and has an atomic number 1. Hence, it has one valence electron.
Similarly, oxygen (O) belongs to Group 16 (the chalcogen family) and has an atomic number of 8.
Hence, it has six valence electrons. The other two electrons are under heavy attraction from the nucleus and cannot participate in chemical bonding.
The H2O2 molecule consists of two H atoms and two O atoms.
Thus, the total number of valence electrons (n1) =
2 x (Number of valence electrons in H) + 2 x (Number of Valence electrons in O)
= 2 x 1 + 2 x 6 = 14
Step 2: Calculate the number of electrons needed to complete the octet of each atom in the molecule.
Note that the H atom needs a duet instead of an octet to achieve the stable, inert electronic configuration of Helium.
Thus, the total number of electrons needed to achieve inert configuration (n2) =
2 x 2 (for H) + 2 x 8 (for O)
= 4 + 16 = 20
Step 3: Calculate the number of bonding electrons.
Number of bonding electrons (n3) = n2 – n1 = 20 – 14 = 6
Step 4: Calculate the number of bond pairs.
Number of bond pairs = n3/2 = 6/2 = 3
Step 5: Calculate the number of non-bonding electrons.
Number of non-bonding electrons (n4) = Valence Electrons – Bonding Electrons
= n1 – n3
= 14 – 6 = 8
Step 6: Calculate the number of lone pairs.
Number of lone pairs = n4/2 = 8/2 = 4
Step 7: Draw the skeletal structure of H2O2 without the electron pairs.
Generally, the least electronegative atom is chosen as the central atom in a lewis structure.
However, in H2O2, there cannot be a central atom as there are two atoms of each kind. The matter is further worsened because the least electronegative element, H, is a monovalent element, i.e., it can form only one bond.
But, from our previous calculation, we can guess the skeletal structure.
We will later verify our guess by using the concept of formal charges. We draw the two O atoms in the center and the two H atoms on the edges for the guess.
Step 8: Place the valence electrons in the skeletal structure.
We have three bond pairs of electrons (colored in blue). We place them to create two oxygen-hydrogen bonds and one oxygen-oxygen bond.
The H atoms have achieved the inert configuration as they have 2 electrons each. But, each O atom is falling short by 4 electrons to achieve an octet. But we have four lone pairs too (colored in red).
We can place them to fulfill the octet of O atoms.
Thus, we can see that the O atoms have a complete octet and the H atoms have a complete duet of electrons.
This inert electronic configuration is vital to ensure the stability of the molecule.
Step 9: Calculate the Formal Charge on all atoms.
We can check the validity of our Lewis structure by using the concept of formal charge. Formal charge on an atom is defined as follows:
Formal Charge = Valence Electrons – (0.5 x Bonding Electrons) – Non-Bonding Electrons
The number of bonding and non-bonding electrons can be found from the Lewis structure.
Formal Charge for each O atom = 6 – (0.5 x 4) – 4 = 6 – 6 = 0
Formal Charge for each H atom = 1 – (0.5 x 2) – 0 = 1 – 1 = 0
Total Charge on the molecule = Sum of formal charges on atoms = 0 + 0 + 0 + 0 = 0
This is consistent as H2O2 is indeed an uncharged neutral molecule. Thus, our Lewis structure is correct.
Molecular Geometry of H2O2
We employ Valence Shell Electron Pair Repulsion (VSEPR) to determine the molecular geometry of H2O2.
VSEPR theory states that the molecules adopt a geometry to minimize the repulsion between the electron clouds.
These interactions are of three kinds. In descending order of strength, they are as follows:
1. Lone pair – Lone pair repulsion
2. Bond pair – Bond pair repulsion
3. Bond pair – Bond pair repulsion
We shall use the VSEPR geometry table to determine the geometry.
Here, A is the central atom, X is the substituent, and E is the electron lone pair.
The table is valid only for molecules with a central atom (A). However, we can approximate a part of the H2O2 molecule as an AX2E2 type molecule.
We write H-O as R. Thus, the formula for H2O2 becomes R-O-H. Now, the VSEPR table can be applied. The molecule is of AX2E2 type with two bond pairs and two lone pairs.
Thus, the geometry will be of bent type. But, we have applied the VSEPR model to only one of the O atoms. Having bent geometry at each O atom results in an “open book” structure for the H2O2 molecule.
Note: The bond lengths and angles differ slightly in the solid crystalline phase and the figure shown below (the gas phase). The reasons for this are out of scope for this article. However, the “open book” structure remains the same.
So, far we have dealt with chemical bonding using only the valence electrons.
For a more precise understanding of molecular geometry, we need to know the electronic configuration of the atomic orbitals.
The atomic orbitals are one-electron wavefunctions, and the square of their amplitude gives us the probability of finding the electron in space.
The electronic configuration can be determined using Aufbau’s principle, Hund’s Rule of maximum multiplicity, and Pauli’s exclusion principle.
Electronic Configuration of Oxygen (O): [He] 2s2 2p4
Electronic Configuration of Hydrogen (H): 1s1
Note that the six valence electrons of oxygen lie in s and p orbitals. But, the chemical bonds formed by O have to be energetically equivalent. This is where hybridization steps in.
One s orbital and three p orbitals mix to give rise to four energetically equivalent sp3 orbitals.
We can also calculate the hybridization of O2 using the Steric Number.
Steric Number of O = Number of Atoms bonded to O + Number of lone pairs on O
= 2 + 2 = 4
Steric Number = 4 corresponds to sp3 hybridization.
Two of the four sp3 hybrid orbitals contain the two lone pairs of electrons, while the two sp3 orbitals are available for chemical bonding.
One of the sp3 orbitals overlaps with the 1s of an H atom to form an oxygen-hydrogen covalent bond, while the other sp3 orbital overlaps with the sp3 orbital of another oxygen atom to form an oxygen-oxygen single bond.
The sp3 orbitals are directed towards the corners of a tetrahedron, but the presence of two lone pairs forces the oxygen atoms to adopt a bent shape like the H2O molecule.
Having a bent geometry for both Oxygen atoms results in an open book structure.
Electronegativity is the tendency of an atom to pull the shared pair of electrons towards itself.
Elements with a greater tendency to do so are called electronegative elements, and those with a lesser tendency to do so are called electropositive elements.
Electronegativity has been quantified using the Pauling Scale of electronegativity. A greater number on the Pauling scale corresponds to a more electronegative element.
Oxygen (3.44) is more electronegative than hydrogen (2.20). Hence, the O atoms will have a partial negative charge δ- while the H atoms will have a partial positive charge δ+.
The chemical dipole moment vector is defined as
μ = δ x d
where δ is the partial charge and d is the distance between the atoms.
By convention, it points from the electropositive molecule to the electronegative molecule. Molecules with a permanent dipole moment are known as polar molecules.
By the values of Pauling electronegativity, we can infer the direction of the dipole moment vectors (colored in grey) in H2O2.
The O-O bond is nonpolar as there is no electronegativity difference between the two O atoms.
The dipole moment vectors are equal in magnitude, but they are not antiparallel to each other. Therefore, the resultant dipole moment vector for the molecule is non-zero.
Hence, H2O2 is a polar molecule.
I have also published an article specifically on the polarity of H2O2.
Here, in this article, we have described the hydrogen peroxide molecule H2O2.
We have used Lewis Structures and VSEPR theory to predict the molecule’s molecular geometry, hybridization, and polarity.
The H2O2 molecule exhibits the sp3 hybridization with an open book-like structure. It is a polar molecule due to its bent shape.
Let us know if you have any questions floating in your mind.