[NO2]– is the chemical formula for the nitrite ion. It is a polyatomic ion having a molar mass of 46.005 g/mol. The nitrite ion is used in food processing as it inhibits bacterial growth. Medically, it is also used as a vasodilator to relieve cardiac pain.
To find out whether the nitrite [NO2]– ion is polar or non-polar, continue reading the article.
Is [NO2]– polar or non-polar?
Nitrite [NO2]– is a polar molecular ion. Itconsists of a nitrogen (N) atom and two oxygen (O) atoms. The nitrogen atom is present at the center of the molecular ion, having a lone pair of electrons. In contrast, two oxygen (O) atoms occupy terminal positions, one on each side of the central N-atom, thus forming a bent shape.
An electronegativity difference of 0.40 units exists between a nitrogen and an oxygen atom in each of the N=O and N-O bonds in the NO2– ion. Thus, both covalent bonds are individually polar in the NO2– molecular ion and possess a specific dipole moment value (symbol µ).
It is due to the bent shape of NO2– that the dipole moments of N-O and N=O bonds do not get canceled in the ion overall. As a result, the electron cloud stays non-uniformly spread thus, NO2– is overall polar (net µ > 0).
A molecule is polar if there is a non-uniform charge distribution present in it. If the charge distribution gets equally balanced in different parts, then that molecule or molecular ion is considered non-polar.
The following three factors mainly influence the polarity of a molecule or molecular ion:
The electronegativity difference between two or more covalently bonded atoms
Dipole moment
Molecular geometry or shape
Now, let us discuss the effect of the above three factors one by one to prove that the nitrite (NO2–) ion is overall polar.
Factors affecting the polarity of [NO2]–
Electronegativity
It is defined as the ability of an atom to attract a shared pair of electrons from a covalent chemical bond.
Electronegativity decreases down the group in the Periodic Table of elements while it increases across a period.
Greater the electronegativity difference between bonded atoms in a molecule, the higher the bond polarity.
Nitrogen (N) is present in Group V A (or 15) of the Periodic Table. The electronic configuration of nitrogen is 1s2 2s2 2p3. As per this electronic configuration, an N-atom has a total of 5 valence electrons. It is thus short of 3 valence electrons that are required so that a nitrogen atom can achieve a complete octet electronic configuration.
Conversely, oxygen (O) belongs to Group VI A (or 16) of the Periodic Table. The electronic configuration of an oxygen atom is 1s2 2s2 2p4. So each O-atom has a deficiency of 2 more valence electrons for it to complete its octet.
Hence in NO2–, two O-atoms are covalently bonded to the central N-atom, one on each side of the bent or angular V-shape. One O-atom is bonded via a double-covalent bond, while the other O-atom is single-covalently bonded to the central N-atom in NO2–.
In this way, both bonded oxygen atoms attain a complete octet configuration via covalent chemical bonding.
3 valence electrons consumed in chemical bonding out of the 5 initially present leaves behind 2 valence electrons, i.e., 1 lone pair on the central N-atom in NO2–. This lone pair of electrons on the central N-atom leads to distortion in NO2– shape, geometry, and symmetry.
Atom
Electronic configuration
Valence electrons
Nitrogen (7N)
1s22s22p3
5
Oxygen (8O)
1s22s22p4
6
Oxygen (E.N = 3.44) is more electronegative than nitrogen (E.N = 3.04). There is an electronegativity difference of 0.40 units between an N and an O-atom in each N-O single bond and N=O double bond present in NO2–. The more electronegative O-atom strongly attracts the shared electron cloud away from the central N-atom in each covalent bond.
Thus, the nitrogen atom present at the center of NO2– gains a partial positive (Nδ+) charge, while each terminal oxygen atom obtains a partial negative charge (Oδ-). As a result, both N-O and N=O bonds are individually polar in the nitrite (NO2–) ion, as shown below.
Dipole Moment
The dipole moment is the product of electrical charge (Q) and bond length (r) between two bonded atoms. It is a vector quantity expressed in Debye (D) units.
It is represented by a Greek symbol µ and measures the polarity of a bond.
The dipole moment of a polar covalent bond conventionally points from the positive center to the center of the negative charge.
So in NO2–, the dipole moment of both N=O and N-O polar bonds points from Nδ+ to Oδ- (as shown below).
Molecular geometry
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, NO2– is an AX2E1-type molecular ion. To one N-atom at the center (A), 2 bond pairs (X) are attached, and there is one lone pair (E) on the central nitrogen atom.
So, the molecular geometry or shape of NO2– is bent, also known as angular or V-shaped. The O=N-O bond angle is 134° in NO2–, different from the ideal bond angle in an AX3-type trigonal planar molecule, i.e., 120°.
The presence of one lone pair of electrons on the central nitrogen atom in [NO2]– ion leads to lone pair-bond pair electronic repulsions along with a bond pair-bond pair repulsive effect.
This strong repulsive effect pushes the bonded atoms away from the lone pair at the center. Thus, NO2– adopts an asymmetric bent shape.
It is due to this asymmetric bent shape that the dipole moments of N-O and N=O polar bonds do not get canceled in NO2–. The electron cloud is not uniformly distributed over the molecular ion. Thus, NO2– is overall polar (net µ > 0).
Difference between polar and nonpolar?
Polar molecule
Non-polar molecule
Atoms must have a difference in electronegativity
Atoms may have the same or different electronegativity values
Unequal charge distribution overall
Equal charge distribution overall
Net dipole moment greater than zero
Net dipole moment equals to zero
Examples include water (H2O), ethanol (CH3CH2OH), ammonia (NH3), sulfur dioxide (SO2), bromine trifluoride (BrF3), nitrite [NO2]– ion, etc.
Examples include oxygen (O2), nitrogen (N2), methane (CH4), nitrate [NO3]– ion, etc.
NO2– has polar bonds present due to a specific electronegativity difference of 0.40 units between the bonded nitrogen and oxygen atoms in both the N-O and N=O bonds.
The presence of one lone pair of electrons on the central nitrogen atom in NO2– leads to an asymmetrical bent or V-shape of NO2–. The N-O and N=O dipole moments do not get canceled.
Thus, NO2– is overall polar with a net dipole moment µ > 0.
What are the formal charges present on bonded atoms in NO2–?
Formal charge on an atom = [ valence electrons – non-bonding electrons- ½ (bonding electrons)]
In NO2–, zero formal charges are present on the central N-atom and on the double-bonded oxygen atom, while a -1 formal charge is present on the single-bonded O-atom; thus, the overall charge present on the nitrite ion is 0 + (0) + (-1) = -1.
Why [NO2]– ion is polar, but [NO3]– ion is non-polar?
In the nitrite [NO2]– ion, there is a lone pair of electrons on the central N-atom. Electronic repulsions distort the shape and geometry of the molecular ion. An asymmetric bent shape ensures the N-O and N=O dipole moments do not get canceled. Thus it is polar (net µ > 0).
Contrarily, in the nitrate [NO3]– ion, there is no lone pair of electrons on the central N-atom. The ion thus adopts a perfectly symmetrical trigonal planar shape and geometry. The dipole moments of polar N-O and N=O bonds get canceled equally. It is thus non-polar (net µ =0).
How do you compare the polarity of NO2– with NO2?
Both the nitrite ion (NO2–) and nitrogen dioxide (NO2), a neutral molecule, are polar in nature.
A specific electronegativity difference between respective nitrogen and oxygen atoms accounts for bond polarity, while the distortion present in the asymmetric bent shape of NO2– and NO2 accounts for an overall molecular polarity.
How do you compare the polarity of NO2– with NO2+?
NO2– is polar, while NO2+ (nitronium ion) is non-polar. Both are made up of individually polar N-O bonds.
However, in NO2– the central N-atom has a lone pair of electrons which leads to distortion in its shape and geometry. It is due to the bent shape of NO2– that N-O dipole moments do not get canceled, so it is polar (net µ > 0).
Contrarily, there is no lone pair of electrons on the central N-atom in NO2+; thus, it has a symmetrical linear shape. The N-O dipole moments get canceled to yield an overall non-polar molecular ion (net µ = 0).
Summary
Nitrite [NO2]– is a polar molecular ion.
It consists of one N-O and one N=O polar bond having an electronegativity difference of 0.4 units between an oxygen and a nitrogen atom.
Nitrite [NO2]– has a bent or V-shape with a 134° bond angle.
The electron cloud is not uniformly distributed in the asymmetric NO2–.
The dipole moment of the N=O double bond does not get canceled with the dipole moment of the N-O single bond. Thus the charged electron cloud stays non-uniformly distributed over the molecular ion.
The net dipole moment of NO2– is > 0, so it is overall polar.
Vishal Goyal is the founder of Topblogtenz, a comprehensive resource for students seeking guidance and support in their chemistry studies. He holds a degree in B.Tech (Chemical Engineering) and has four years of experience as a chemistry tutor. The team at Topblogtenz includes experts like experienced researchers, professors, and educators, with the goal of making complex subjects like chemistry accessible and understandable for all. A passion for sharing knowledge and a love for chemistry and science drives the team behind the website. Let's connect through LinkedIn: https://www.linkedin.com/in/vishal-goyal-2926a122b/
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