A radical is known to consist of an unpaired valence electron and therefore tends to be highly chemically reactive. This property makes this radical short-lived (unless treated with certain organic molecules). It has a molar mass of around 16.0226 g/mol and is the neutral form of amide ions or azanide ions. Here, in this article, we will focus on the azanide anion (NH2-). NH2- is the conjugate acid of hydridonitrate (H-N2-) and conjugate base of ammonia (NH3). 2Li + 2NH3 —–> 2LiNH2 + H2 The above reaction shows how metal amides like Lithium Amide are being produced from liquid ammonia solution and Li metal. Azanide is a monovalent and inorganic anion having a negative charge of -1. NH2- is not stable as such and therefore it is found to exist as a hydrazine compound. Amidyl group( -NH-) and azanidylidene group (=N-) are substituent groups from azanide. ( Amidyl group is formed due to a loss of a proton from NH2-). Let us now get into our topic of discussion and talk about the chemical bonding nature of NH2-.
NH2 Lewis Structure
The most initial and important step towards finding out the bonding nature of any chemical molecular structure is to draw its Lewis Structure. Lewis Structure is a 2D diagram to represent the internal bonding between constituent atoms in a molecule or ion. Here, we have three key concepts to understand before we can proceed to sketch the Lewis Structure of any given molecule, in this case, azanide ion.
- Valence electrons and electron dot notations: We already know what valency is. It is defined to be the combining capacity of an atom to form bond structures. When two atoms share electron pairs, a bond is formed. Accordingly, there can be single, double, and triple bonds. Valence electrons refer to the electrons in the outermost shell of an atomic nucleus. We denote these as dots while drawing the Lewis Structure, also known as electron-dot structure. The bonds are denoted by straight lines.
- Octet rule If we have a look into the modern periodic table, we can see that the main group elements are present between groups 1-17. These elements have a tendency to attain the outer shell electronic configuration of noble gas elements, i.e. they tend to achieve a valency of 8 while bond formation. This is known as the octet rule or octet fulfillment. However, there are certain exceptions to this rule.
- Formal Charge If we assume that electrons will always be shared in equal proportions between atoms while forming a molecular structure, then the individual electric charge assigned to the atoms is said to be the formal charge. While finding out the pictorial representation of the molecule, it is a crucial step to check whether the atoms are in their least possible formal charge values.
Let us start drawing lewis structure for NH2
Nitrogen belongs to period 15 and has 5 valence electrons whereas hydrogen belongs to period 1 and has only 1 valence electron. Other than this, we have an electron that gives us the negative charge of -1. The total number of valence electrons in an NH2- anion = 5 + 2*1 + 1 = 8. If we compare electronegativities, Hydrogen has less value than Nitrogen. The atom having less electronegativity value forms the central atom in general. However, hydrogen, having only one electron and not being capable of producing several covalent bonds, do not remain or act as the central atom. So, nitrogen forms the central atom in this case.
We have placed the atoms with nitrogen in the center.
We have placed all the valence electrons and denoted them by dot notations. As we can check, the nitrogen atom has achieved octet configuration. As for hydrogen, as an exception to the general octet rule, it follows helium configuration and both the atoms have achieved so.
Now, we have to check the formal charge values of all the constituent atoms.
Formal charge of N= 5 – 0.54 – 4 = -1. Formal charge of H= 1 – 0.52 – 0 = 0. As we can see, there is a net negative charge on the NH2 structure.
NH2 Molecular Geometry
Do you know that we can predict the 3D molecular shape of a molecule from its Lewis Structure? For this, we need to understand the concept of VSEPR or Valence Shell Electron Pair Repulsion theory. This theory tells us that since electrons are like charges, the negative cloud atmosphere formed by them surrounding each atomic nucleus tends to experience repulsion. We need to minimize the repulsion occurring between electrons to stabilize a molecular structure. Let us check the VSEPR notation for NH2-. AXnEx notation: A is the central atom, X = surrounding atoms or electron bonded pairs of the central atom, E= lone pair on the central atom. For azanide anion, the notation is AX2E2.
As we can see in the above VSEPR chart, we have a bent molecular geometry for NH2-. It has a bond angle of 104.50, much less than the general ideal 109.50 value. This occurs due to the strong repulsive power of the two lone pairs on the central N atom.
Electron Geometry
If we look at the Lewis Structure of NH2- ion, we can see that we have four electron-rich regions around our central nitrogen atom. There are two bonded electron pairs around N to form single covalent bonds each with a Hydrogen atom on either side. Apart from this, N has two unbonded or lone electron pairs. These four electron pairs (bonded and unbonded) result in a tetrahedral electron geometry.
NH2 Hybridization
Orbital hybridization is a noteworthy topic in chemistry. It is a mathematical model to explain the phenomenon of covalent bonding. If we consider electrons to behave as both particles and waves, we can say that we cannot predict the exact location of electrons around an atomic nucleus as stated by Heisenberg’s Uncertainty Principle. Hence, we talk about orbitals which give us an idea of the probability of the presence of electrons in a regional space. Atomic orbitals of the same atom of equivalent energies come together and fuse to form hybridized orbitals according to Valence Bond Theory (VBT). For example, an ‘s’ and a ‘p’ orbital come together to form an sp hybrid orbital, one ‘s’ and two ‘p’ orbitals ( e.g. px, py ) fuse to form an sp2 hybrid orbital, and so on. Let us check what the hybridization of NH2- ion is: For finding out hybridization, we can use the following formula: H= 1/2 ( V + M – C + A) H= hybridization type V= number of valence electrons M= monovalent atoms C= cationic charge A= anionic charge Here, V=5, M=2. C=0, A=1. H = ½( 5 + 2 – 0 + 1) = 4. For Nitrogen atom, no. of valence electrons= 5. If we look at the electronic configuration of N, we can see: N: 1s2 2s2 2p3 The 1s orbital, not being an outer shell orbital does not take part in the hybridization. We only take into account the valence shell electrons. So, the s and three 2p orbitals (px, py, pz) fuse to form sp3 hybrid orbitals.
NH2 Polarity
Polarity is another important concept of chemistry that we will discuss in this article with regard to NH2- ion. Polarity is the concept of charge separation inside a molecule. To decipher the polar nature of azanide, we need to look into the Pauling Electronegativity chart.
The knowledge of electronegativity i.e the power or degree to which an atom can attain negatively charged electrons required to find out whether a given molecule is polar or non-polar. As we can find out, N has an electronegativity value of 3.04 and H has an electronegativity value of 2.20. So, there is a huge electronegativity difference in the N-H bonds which results in induction of dipole moment. It is calculated as the product of the charge and the distance between both the charges. There will be partial positive and negative charges (δ+ and δ-) at both ends. Hence, polar bonds are formed as a result. For core reasons of the polarity of NH2, you must check out the polarity of NH2. If we look at the 3D geometry of the molecule, we can see that it is not linear due to the presence of repulsive forces via lone pairs as well as bond pairs. This also contributes to the polar nature of NH2- ion.
Conclusion
Here, in this article, we have explained the bonding nature of azanide or amide ion. We have included in our discussion the Lewis Structure, molecular geometry, Hybridization, and Polarity. Happy learning chemistry!
