NO2 (Nitrogen Dioxide) Lewis Dot Structure - Trending Science (2023)

Nitrogen Dioxide (NO₂) is a covalent compound consisting of a central nitrogen atom singly bonded to one oxygen atom and doubly bonded to another oxygen atom. At room temperature, nitrogen dioxide is a red-brown gas with a density of 1.8 g/dm.³🇧🇷 it isslightly toxic to humans, due to its tendency to react in the human body and produce reactive nitrogen and oxygen species that can damage internal structures.

"A small bubble of air remained unabsorbed...if there is a portion of the phlogistic air [nitrogen] in our atmosphere that is distinct from the rest and cannot be reduced to nitrous acid, we can safely conclude that it is no more than 1/120." of the whole." – Henry Cavendish

Nitrogen dioxide does not have a single Lewis structure due to its relatively bizarre electron configuration. The position of the double bond changes over time, which means that any of the oxygen atoms can form a double bond with the nitrogen atom at any time. As such, nitrogen dioxide is represented by the Lewis resonance structure:

Nitrogen dioxide requires a Lewis resonance structure because itelectronic configurationconstantly oscillating between the two forms. The "true" electron configuration of nitrogen dioxide is considered to be the average of the two resonance structures given above. The Lewis structure of nitrogen dioxide is also interesting because there is a single unpaired valence electron at the central nitrogen atom. Compounds with unpaired electrons are sometimes referred to as "free radicals". This unpaired electron explains nitrogen dioxide's reactive behavior as it has a strong desire to fill this open space with electrons.

Let's step back and review the rules for drawing a Lewis structure. We'll go step by step to see how we can construct a Lewis structure for most main group compounds, including nitrogen dioxide.

Lewis Structures: Basic Concepts

Simply put, a Lewis structure is a pictorial representation of atomic structure andElectronicConfiguration of an atom or acomposed🇧🇷 Individual atoms are represented by their unique chemical symbol, electrons are represented as single dots, and shared pairs of electrons are represented by a single dash (-) for a lone pair, a double slash (=) for a double pair, and a triple. Forward slash (≡) for a triple pair.

“Any chemical substance, whether natural or artificial, falls into one of two main categories depending on the spatial character of its form. A distinction is made between substances that have a plane of symmetry and those that do not. The former belong to the mineral, the latter to the living world." -Louis Pasteur

The purpose of a Lewis structure is to see how electrons are arranged in an atom or compound. Lewis structures are based on the octet rule: the empirical observation that atoms tend to form bonds until they have a full valence shell of 8 electrons. The only exception to the octet rule is hydrogen, which will only bond until it has 2 valence electrons.

Valence electrons are represented as pairs of dots., where each point represents a single electron. Atoms form covalent bonds by sharing their valence electrons with other atoms. For example, a single chlorine atom has 7 valence electrons; 3 pairs and a free electron. Two chlorine atoms share their unpaired electrons, so each atom has a full octet of electrons, forming a chlorine molecule (Cl₂🇧🇷 In general, it is like thiscovalent bondswork. Atoms share valence electrons until each atom has a full octet. When all the valence electrons are paired but an atom does not yet have a full octet, the electron pairs move to form double and triple bonds. The total number of electrons in a Lewis structure is equal to the sum of the number of valence electrons in each atom.

Lewis structures tell you about the atomic arrangement and distribution of electrons in an atom or compound. Although Lewis structures alone do not provide explicit information about a molecule's three-dimensional geometry, the rules for writing Lewis structures can be combined with the rules governing molecular geometry to predict the shape of a compound.

Rules for making Lewis structures

Let's repeat the rules for making the Lewis structure using nitrogen dioxide as our test example.

Step 1. Determine the total number of valence electrons.

The first step is to figure out how many electrons the diagram should have. The total number of electrons in a Lewis structure must equal the sum of the valence electrons in each atom. The number of valence electrons in an element can be determined by looking at its group number in theperiodic table🇧🇷 In general, elements in groups 1 and 2 have 1 and 2 valence electrons, respectively. Group 13-18 elements have 3, 4, 5, 6, 7, and 8 valence electrons, respectively. Group 3-12 elements are transition metals that can have different numbers of valence electrons.

In our case, nitrogen dioxide consists of 1 nitrogen atom and two oxygen atoms. Nitrogen is a group 15 element and therefore has 5 valence electrons, while group 16 oxygen has 6 valence electrons. There are two oxygen atoms, so the total number of valence electrons in our diagram is:

5(1) + 6(2) =17 electrons

Our diagram should have a total of 17 electrons.

Step 2: Draw the atomic structure of the compound

Now that we have the number of valence electrons, we can start building the diagram. When the compound is diatomic (two atoms), the structure is simple; Atoms can be placed next to each other. In compounds with three or more atoms, there is usually a central atom that shares multiple bonds with terminal atoms. Generally, in triatomic or larger compounds, the central atom is leastelectronegativeElement.

In our case we have a triatomic compound, so our structure will likely have a central atom bonded to multiple end atoms. Nitrogen is less electronegative than oxygen (3.04 < 3.44), so nitrogen is our center. The placement of the symbols gives us:

Step 3: Pair electrons so that each atom has at least one single bond

Next, we can start incorporating electrons into our model. First we go through and place a single bond between each atom. Each individual bond has 2 electrons, so we subtract those electrons from the total to figure out how many more we have to put in.

In our case we put two single bonds, one between each atom, which looks like this:

If we put 2 pairs, we put 4 electrons in total. Now we have 17-4 =13to place more electrons.

"Modern chemistry, with its strong generalizations and hypotheses, is a good example of how far beyond the limits of the senses the human mind can go in exploring the unknown." - Horace G. Deming

Step 4: Pair electrons starting with the terminal atoms until each has a full octet.

Then we put the remaining electrons. Start with the terminal atoms and fill in the dots until each atom has a total of 8 valence electrons. If you have extra atoms, place them in pairs or as single electrons on the central atom.

Starting with the terminal oxygen atoms, we first wrap 6 electrons around each so they have a full octet. We put all the remaining electrons on the nitrogen atom. Adding 6 electrons to each oxygen atom gives 12, so we put the only remaining electron on the nitrogen atom:

After putting in those 13 electrons, we now have 13-13 =0 electronsto put left. But we're not done yet, because our central atom doesn't have a full octet yet, nitrogen currently only has 5 electrons; 2 pairs and a single unpaired electron.

Step 5: Move the pairs of electrons to form double and triple bonds until each atom is an octet or as close to an octet as possible.

When all the electrons have been placed and some atoms still don't have a full octet, the links form double and triple bonds to ensure each atom gets as close to 8 electrons as possible. Just move the electron pairs from the end atoms to form double and triple bonds.

In our case, we have all of our electrons bound, but nitrogen only has 5 valence electrons. The movement of a lone pair of electrons from one of the oxygen atoms creates a double bond with nitrogen, creating 7 electrons. Moving more electron pairs would give nitrogen more than 8 electrons, so we've gone as far as we can and our Lewis structure should look like this:

resonance structures

In the last step of drawing our Lewis diagram, we had to choose a pair of electrons that should move to form a double bond. We chose the oxygen atom on the left, but we couldn't have chosen the oxygen atom on the right to get something like this:

The answer is yes, this is also a valid Lewis structure for nitrogen dioxide. However, this structure is obviously different from the previous one; the double bond is on the right instead of on the left. If both Lewis structures are legitimate, then what is the "true" Lewis structure of nitrogen dioxide? The answer is:both.

In cases where more than one legitimate Lewis structure exists for a compound, the full Lewis structure is presented as the average of the multiple structures. These structures are known asresonance structuresand are used for compounds whose electronic configuration cannot be fully represented by a single Lewis diagram. A resonance structure for our two nitrogen dioxide diagrams looks like this:

The "real" structure ofnitrogen dioxideis interpreted as a combination of the two diagrams. Resonance structures are possible because in some compounds the electron pairs arerelocatedand oscillate between one configuration and another. Resonance structures are necessary because the atomic configurations of some molecules cannot be accurately captured with a single Lewis structure.

Limitations of Lewis Structures

Following the rules of the Lewis structure should allow you to construct a Lewis structure for most compounds made up of main group elements in the s and p blocks of the periodic table. Some families of elements do not always obey the rules for making Lewis structures. crossingrails, for example, often do not follow the octet rule and can bind and gain up to 12 valence electrons. The further down one goes in the periodic table, the less the basic rules for Lewis structures apply due to the extremely heavy nuclei and strong electromagnetic properties of large elements.