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How to count formal charge in NO 2 - ?

Formal charge:- formal charge of a compound can be calculated by the formula: fc = v − n − b 2 here, fc= formal charge, v= number of valence electrons in a free atom, n= total number of non-bonding electrons, b= total number of bonding electrons formal charge of no 2 - :- structure of no 2 - is: step 1: formal charge of nitrogen here nitrogen is the free atom and the number of valence electrons of it is 5. number of non-bonding electrons is 2 and bonding electrons are 6 ∴ formal charge of nitrogen is fc = v − n − b 2 ⇒ fc = 5 - 2 - ( 6 2 ) ⇒ fc = 5 - 5 ⇒ fc = 0 step 2: formal charge of double-bonded oxygen here oxygen is the free atom and the number of valence electrons of it is 6. number of non-bonding electrons is 4 and bonding electrons are 4 ∴ formal charge of double bonded oxygen is fc = v − n − b 2 ⇒ fc = 6 - 4 - ( 4 2 ) ⇒ fc = 2 - 2 ⇒ fc = 0 step 3: formal charge of single-bonded oxygen here oxygen is the free atom and the number of valence electrons of it is 6. number of non-bonding electrons is 6 and bonding electrons are 2 ∴ formal charge of single-bonded oxygen is fc = v − n − b 2 ⇒ fc = 6 - 6 - ( 2 2 ) ⇒ fc = 0 - 1 ⇒ fc = - 1 step 4: formal charge of no 2 - adding all the three different formal charges of nitrogen and oxygens we get the formal charge of no 2 - ∴ formal charge of no 2 - is fc = 0 + 0 + ( - 1 ) ⇒ fc = - 1 therefore, the formal charge of no 2 - is - 1.

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How to count formal charge in NO2?

How to count formal charge in nitrite ion.

calculate formal charge of no2

And so #""^(-)O-stackrel(ddot)N=O# ...around each atom from your left to right there are 9, 7, and 8 electrons respectively leading to formal charges of #-1# , #0# , and #0# .

Of course we can distribute the negative charge over the two oxygen centres by resonance , and so #/_O-N-O# #<# #120^@# , i.e. the nitrogen lone pair, which is closer to the nitrogen atom exerts a disproportionate influence on the bond angle.

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calculate formal charge of no2

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NO2 Lewis Structure: Drawings, Hybridization, Shape, Charges, Pairs

Nitrogen dioxide (NO2) is a chemical compound composed of nitrogen and oxygen atoms . It is a reddish-brown gas with a pungent odor and is commonly found in urban areas as a result of air pollution . Understanding the Lewis structure of NO2 is important in determining its chemical properties and reactivity. The Lewis structure provides a visual representation of the arrangement of atoms and electrons in a molecule, helping us understand how the molecule interacts with other substances . In this article , we will explore the NO2 Lewis structure in detail, discussing its electron arrangement , bond formation , and overall molecular shape . So, let’s dive in and unravel the mysteries of NO2!

Key Takeaways

  • The NO2 Lewis structure consists of a nitrogen atom bonded to two oxygen atoms.
  • The nitrogen atom has a l one pair of electrons, while the oxygen atoms have three l one pair s each.
  • The nitrogen-oxygen bonds are represented by single bonds, and the nitrogen -oxygen double bond is represented by a double bond.
  • The formal charge s on the atom s in the NO2 Lewis structure are: nitrogen (-1), one oxygen (+1), and the other oxygen (0).
  • The NO2 molecule has a bent shape due to the repulsion between the l one pair s of electrons on the nitrogen and oxygen atoms .

NO2 Lewis Structure

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Drawing NO2 Lewis structure

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The Lewis structure is a way to represent the bonding and electron distribution in a molecule. In the case of NO2, which stands for nitrogen dioxide, we can draw its Lewis structure to understand its molecular geometry and electron arrangement .

To draw the Lewis structure of NO2, we need to follow a few steps :

Determine the total number of valence electrons: In NO2, nitrogen (N) is in Group 5A of the periodic table , so it has 5 valence electrons. Oxygen (O) is in Group 6A, so each oxygen atom has 6 valence electrons. Since there are two oxygen atoms in NO2, we have a total of 5 + 2(6) = 17 valence electrons .

Identify the central atom: In NO2, nitrogen is the central atom as it is less electronegative than oxygen.

Connect the atom s: Place the nitrogen atom in the center and connect it to the two oxygen atoms using single bonds.

Distribute the remaining electrons: Distribute the remaining electrons around the atom s to satisfy the octet rule . Start by placing l one pair s on the outer atoms (oxygen) and then distribute the remaining electrons on the central atom (nitrogen).

Check for octet rule and formal charge s: Make sure all atoms have an octet of electrons (except hydrogen, which only needs 2 electrons). If necessary, move l one pair s to form double or triple bonds to satisfy the octet rule . Also, check for any formal charge s to minimize their presence .

Explanation of NO2 Lewis structure

The Lewis structure of NO2 shows that nitrogen is bonded to two oxygen atoms. The nitrogen atom has one l one pair of electrons, while each oxygen atom has two l one pair s . The double bond between nitrogen and one oxygen atom is represented by two dots or a double line , indicating the sharing of two pairs of electrons. The single bond between nitrogen and the other oxygen atom is represented by a single dot or a single line , indicating the sharing of one pair of electrons.

The Lewis structure helps us understand the arrangement of electrons in a molecule and predict its molecular geometry . In the case of NO2, the molecule has a bent or V-shaped geometry due to the repulsion between the l one pair s of electrons on the oxygen atoms. This bent shape gives NO2 a polar nature , with the oxygen atoms being slightly negative and the nitrogen atom being slightly positive.

Valence electrons in NO2 Lewis structure

Valence electrons are the electrons in the outermost energy level of an atom. In the Lewis structure of NO2, we consider the valence electrons of nitrogen and oxygen to determine the total number of electrons available for bonding.

Nitrogen, being in Group 5A, has 5 valence electrons. Oxygen, being in Group 6A, has 6 valence electrons. Since there are two oxygen atoms in NO2, we multiply the number of valence electrons of oxygen by 2. Therefore, the total number of valence electrons in NO2 is 5 + 2(6) = 17.

Valence electrons play a crucial role in determining the chemical properties and reactivity of a molecule. They are involved in forming chemical bonds and determining the electron distribution in a molecule.

Octet rule in NO2 Lewis structure

The octet rule states that atoms tend to gain, lose, or share electrons in order to achieve a stable electron configuration with 8 electrons in their outermost energy level . This rule helps us understand the formation of chemical bonds and the stability of molecules.

In the Lewis structure of NO2, we can see that nitrogen has 5 valence electrons and each oxygen atom has 6 valence electrons. By sharing electrons through single and double bonds , nitrogen and oxygen can achieve an octet of electrons in their outermost energy level .

The double bond between nitrogen and one oxygen atom satisfies the octet rule for both nitrogen and oxygen. However, the other oxygen atom only has 7 electrons around it. To satisfy the octet rule , one l one pair from the nitrogen atom is moved to form a double bond with the second oxygen atom . This redistribution of electrons allows all atoms in NO2 to have an octet of electrons, fulfilling the octet rule .

Understanding the NO2 Lewis structure and the application of the octet rule helps us predict the stability and reactivity of molecules. It provides insights into the chemical behavior and properties of compounds like nitrogen dioxide.

Hybridization in NO2 Lewis Structure

The Lewis structure of NO2, also known as nitrogen dioxide, is a representation of the molecule’s bonding and electron arrangement . In order to understand the Lewis structure of NO2, it is important to first grasp the concept of hybridization.

Definition of Hybridization

Hybridization is a concept in chemistry that describes the mixing of atomic orbitals to form new hybrid orbitals . These hybrid orbitals have different shapes and energies compared to the original atomic orbitals . Hybridization occurs when atoms bond together to form molecules.

Hybridization in NO2

In the case of NO2, the central nitrogen atom is bonded to two oxygen atoms. To determine the hybridization of the nitrogen atom in NO2, we need to consider the number of electron groups around it. An electron group can be a l one pair or a bond .

In NO2, there are two oxygen atoms bonded to the nitrogen atom, resulting in two electron groups . Additionally, there is one l one pair of electrons on the nitrogen atom. Therefore, the total number of electron groups around the nitrogen atom is three.

Explanation of sp2 Hybridization in NO2

The hybridization of the nitrogen atom in NO2 is sp2. This means that the nitrogen atom in NO2 undergoes hybridization by mixing one 2s orbital and two 2p orbitals to form three sp2 hybrid orbitals . These sp2 hybrid orbitals are arranged in a trigonal planar geometry around the nitrogen atom.

The three sp2 hybrid orbitals in NO2 are used to form sigm a bond s with the two oxygen atoms and accommodate the l one pair of electrons. The remaining p orbital on the nitrogen atom contains one electron , which is involved in pi bonding with one of the oxygen atoms.

To summarize, the sp2 hybridization in NO2 allows the nitrogen atom to form three sigm a bond s and one pi bond , resulting in a trigonal planar molecular geometry .

In conclusion, understanding the hybridization in the NO2 Lewis structure is crucial for comprehending the molecule’s bonding and shape. The sp2 hybridization of the nitrogen atom in NO2 enables it to form three sigm a bond s and one pi bond , leading to a trigonal planar molecular geometry .

Formal Charges in NO2 Lewis Structure

The formal charge s in a NO2 Lewis structure play a crucial role in understanding the distribution of electrons within the molecule. By assigning formal charge s, we can determine the most stable arrangement of electrons and gain insights into the molecule’s reactivity and properties.

Definition of Formal Charges

Formal charges are hypothetical charges assigned to each atom in a molecule or ion. These charges help us understand the distribution of electrons and determine the stability of different resonance structures . The formal charge of an atom is calculated by comparing the number of valence electrons it should have with the number it actually possesses in the Lewis structure.

Calculation of Formal Charges in NO2

To calculate the formal charge s in a NO2 molecule , we need to follow a step-by-step process:

Determine the total number of valence electrons in the molecule. For NO2, nitrogen (N) contributes 5 valence electrons, and each oxygen (O) contributes 6 valence electrons, giving us a total of 5 + 2(6) = 17 valence electrons .

Assign l one pair s of electrons to each atom. Nitrogen requires 3 l one pair s to complete its octet, while each oxygen atom requires 2 l one pair s.

Connect the atom s using single bonds. In the case of NO2, nitrogen forms a double bond with one oxygen atom, and a single bond with the other oxygen atom.

Distribute the remaining electrons as l one pair s to satisfy the octet rule for each atom. In NO2, the remaining 3 electrons are placed as a l one pair on the nitrogen atom.

Calculate the formal charge for each atom. The formula for formal charge is:

Formal Charge = Valence Electrons – Lone Pair Electrons – 0.5 * Bonding Electrons

For example, for the nitrogen atom in NO2, the formal charge is:

Formal Charge = 5 – 3 – 0.5 * 4 = 0

For each oxygen atom, the formal charge is:

Formal Charge = 6 – 2 – 0.5 * 4 = 0

In the NO2 Lewis structure, the nitrogen atom has a formal charge of 0, while each oxygen atom also has a formal charge of 0. This distribution of formal charge s indicates that the Lewis structure is stable and represents the most favorable arrangement of electrons for NO2.

By analyzing the formal charge s, we can conclude that the nitrogen atom in NO2 does not carry any excess or deficient electrons . Similarly, each oxygen atom also has the appropriate number of electrons to maintain stability.

Understanding the formal charge s in the NO2 Lewis structure provides valuable insights into the molecule’s behavior and reactivity. It helps us predict how NO2 interacts with other molecules and how it participates in chemical reactions .

In the next section , we will explore the molecular geometry and bond angle s in the NO2 molecule, further enhancing our understanding of its structure and properties.

Resonance in NO2 Lewis Structure

Resonance is an important concept in chemistry that helps us understand the behavior of molecules. In the case of the NO2 Lewis structure, resonance plays a significant role in determining its structure and properties.

Definition of Resonance

Resonance refers to the delocalization of electrons within a molecule. It occurs when there are multiple valid Lewis structures that can be drawn for a molecule, and the actual structure is a combination or average of these resonance structures . In other words , resonance structures are different ways of arranging the same atoms , but with different electron distribution s.

When we consider the NO2 molecule, we can draw multiple resonance structures . NO2, also known as nitrogen dioxide, consists of a nitrogen atom bonded to two oxygen atoms. The central nitrogen atom has a l one pair of electrons and forms a double bond with one of the oxygen atoms, while the other oxygen atom is bonded by a single bond.

To represent the resonance in the NO2 Lewis structure, we can draw two resonance structures . In the first resonance structure , the double bond is between the nitrogen and the oxygen on the left , while in the second resonance structure , the double bond is between the nitrogen and the oxygen on the right . These resonance structures can be interconverted by moving the double bond and the l one pair of electrons.

Explanation of Resonating Structures in NO2

The presence of resonance in the NO2 Lewis structure affects the overall structure and properties of the molecule. Due to resonance, the actual structure of NO2 is a hybrid of the two resonance structures . This means that the double bond character is shared between the two oxygen atoms, resulting in a more stable molecule .

The resonance in NO2 also affects the bond lengths and bond angle s within the molecule. In the resonance structures , the nitrogen -oxygen bond lengths are equal, and the nitrogen -oxygen-nitrogen bond angle is approximately 134 degrees. However, in the actual structure, the bond lengths are intermediate between single and double bonds , and the bond angle is slightly less than 134 degrees .

The presence of resonance also influences the polarity of the NO2 molecule. Each resonance structure has a partial positive charge on the nitrogen atom and a partial negative charge on the oxygen atoms. In the actual structure, the polarity is distributed over the molecule, resulting in a polar molecule.

In conclusion, resonance in the NO2 Lewis structure is a fascinating phenomenon that arises due to the delocalization of electrons. It leads to the formation of multiple resonance structures , which contribute to the overall stability , structure, and properties of the NO2 molecule. By understanding resonance , we can gain valuable insights into the behavior of molecules and their chemical reactions .

Bond Angle in NO2 Lewis Structure

The bond angle in the NO2 Lewis structure plays a crucial role in determining the shape and properties of the molecule. Understanding the bond angle is essential for predicting the molecule’s behavior and its interactions with other molecules . In this section , we will explore the definition of bond angle , discuss the bond angle in the NO2 Lewis structure, and explain the bent shape of NO2.

Definition of Bond Angle

The bond angle is the angle formed between two adjacent bonds in a molecule. It is measured in degrees and provides valuable information about the molecular geometry and the arrangement of atoms in a compound . The bond angle is influenced by various factors , including the number of electron pairs around the central atom and the repulsion between these electron pairs .

To understand the bond angle in the NO2 Lewis structure, let’s first take a look at its molecular geometry . NO2, also known as nitrogen dioxide, consists of a central nitrogen atom bonded to two oxygen atoms. The Lewis structure for NO2 shows that nitrogen has one l one pair and two single bonds with oxygen.

In the NO2 Lewis structure, the central nitrogen atom is surrounded by three electron pairs – two from the oxygen atoms and one l one pair . These electron pairs repel each other, causing the molecule to adopt a bent shape. The bond angle in the NO2 Lewis structure is approximately 134 degrees.

Explanation of Bent Shape in NO2

The bent shape of NO2 can be explained by considering the repulsion between electron pairs. The l one pair on the central nitrogen atom exerts a greater repulsive force compared to the bonding pairs. As a result, the bonding pairs are pushed closer together, leading to a decrease in the bond angle .

The repulsion between the l one pair and the bonding pairs causes the NO2 molecule to bend, resulting in a bond angle less than the ideal 180 degrees . This bent shape is also influenced by the electronegativity difference between nitrogen and oxygen, which leads to a polar molecule.

In summary, the bond angle in the NO2 Lewis structure is approximately 134 degrees, indicating a bent shape. This bent shape is a result of the repulsion between the l one pair and the bonding pairs, as well as the electronegativity difference between nitrogen and oxygen. Understanding the bond angle and molecular geometry of NO2 is crucial for comprehending its chemical behavior and interactions with other substances .

Lone Pairs in NO2 Lewis Structure

In the Lewis structure of NO2, l one pair s play a crucial role in determining the molecule’s geometry and properties. Let’s explore the definition of l one pair s, the number of l one pair s in NO2, and the impact they have on the molecule’s geometry .

Definition of Lone Pairs

L one pair s, also known as non-bonding pairs , are pairs of electrons that are not involved in bonding with other atoms . In a Lewis structure , these electrons are represented as dots around the atom . The presence of l one pair s affects the overall shape and polarity of a molecule.

Number of Lone Pairs in NO2

In the NO2 molecule, there are two oxygen atoms bonded to a central nitrogen atom. To determine the number of l one pair s in NO2, we need to consider the valence electrons of each atom. Nitrogen has five valence electrons , while oxygen has six. Therefore, the total number of valence electrons in NO2 is:

(1 × 5) + (2 × 6) = 17

To distribute these electrons , we first form single bonds between the nitrogen atom and each oxygen atom. This accounts for four electrons (two from each oxygen). The remaining 1 3 electrons are then placed as l one pair s around the atom s.

Impact of Lone Pairs on NO2 Geometry

The presence of l one pair s in NO2 affects its geometry and bond angle s . In NO2, the nitrogen atom is surrounded by two oxygen atoms and one l one pair . This arrangement gives rise to a bent or V-shaped molecular geometry .

The repulsion between the l one pair and the bonding pairs causes a distortion in the molecule’s shape . The bond angle between the nitrogen -oxygen bonds is less than the ideal 120 degrees due to this repulsion . In the case of NO2, the bond angle is approximately 134 degrees.

The presence of l one pair s also influences the polarity of the molecule. The electronegativity of oxygen is higher than that of nitrogen, resulting in a polar covalent bond between nitrogen and oxygen. The l one pair on nitrogen further enhances the polarity of the molecule, making NO2 a polar molecule.

To summarize, the NO2 molecule has one l one pair on the central nitrogen atom, which affects its geometry , bond angle , and polarity. The presence of the l one pair causes a bent molecular shape and a bond angle of approximately 134 degrees. Additionally, the l one pair contributes to the overall polarity of the molecule.

In the next section , we will delve into the resonance structure of NO2 and its implications on the molecule’s stability and reactivity.

Polar or Nonpolar Nature of NO2 Lewis Structure

Definition of polar and nonpolar molecules.

Before we delve into the polar or nonpolar nature of the NO2 Lewis structure, let’s first understand what it means for a molecule to be polar or nonpolar.

In chemistry, polarity refers to the distribution of electrons in a molecule. A polar molecule has an uneven distribution of electron density, resulting in a separation of positive and negative charges . On the other hand , a nonpolar molecule has an even distribution of electron density, with no separation of charges.

The polarity of a molecule is determined by the difference in electronegativity between the atom s involved in the chemical bond . Electronegativity is a measure of an atom’s ability to attract electrons towards itself. When two atoms with different electronegativities form a bond , the more electronegative atom pulls the shared electrons closer to itself, creating a polar bond .

Determining Polarity of NO2 Lewis Structure

Now, let’s apply this knowledge to the NO2 Lewis structure to determine its polarity .

The NO2 molecule , also known as nitrogen dioxide, consists of one nitrogen atom (N) and two oxygen atoms (O). To draw the Lewis structure of NO2, we start by counting the total number of valence electrons in the molecule. Nitrogen contributes 5 valence electrons, while each oxygen contributes 6 valence electrons, giving us a total of 5 + 2(6) = 17 valence electrons .

Next, we arrange the atom s in the structure , placing the nitrogen atom in the center and the oxygen atoms on either side . We then connect the atom s using single bonds, which account for 2 electrons each. After connecting the atom s, we distribute the remaining electrons as l one pair s around the atom s to satisfy the octet rule .

In the NO2 Lewis structure, the nitrogen atom is double-bonded to one of the oxygen atoms, while the other oxygen atom has a l one pair . This arrangement gives nitrogen a formal charge of +1 and the oxygen atoms a formal charge of -1 each. The Lewis structure can be represented as follows:

O ╱ N = O ╲ O

Now, let’s analyze the polarity of the NO2 molecule based on its Lewis structure . The nitrogen-oxygen double bond is a polar bond due to the difference in electronegativity between nitrogen and oxygen. Oxygen is more electronegative than nitrogen, so it pulls the shared electrons closer to itself, creating a partial negative charge on the oxygen atom and a partial positive charge on the nitrogen atom.

Additionally, the l one pair on the oxygen atom also contributes to the polarity of the molecule. The presence of the l one pair creates an uneven distribution of electron density, further enhancing the polarity of the NO2 molecule.

Therefore, based on the arrangement of atoms and the polarity of the bonds and l one pair s, we can conclude that the NO2 molecule is polar in nature.

In summary, the NO2 Lewis structure exhib its polarity due to the polar nitrogen-oxygen double bond and the presence of a l one pair on one of the oxygen atoms. Understanding the polarity of molecules is crucial in various chemical reactions and interactions, as it influences the behavior and properties of substances.

VSEPR Model and NO2 Lewis Structure

The VSEPR ( Valence Shell Electron Pair Repulsion) model is a useful tool in predicting the shape and geometry of molecules. By considering the repulsion between electron pairs, we can determine the arrangement of atoms in a molecule. In this section , we will explore the application of the VSEPR model to the NO2 Lewis structure and discuss the electron geometry of NO2.

Overview of VSEPR Model

The VSEPR model is based on the principle that electron pairs in the valence shell of an atom repel each other. This repulsion leads to a specific arrangement of atoms in a molecule, which determines its shape and geometry. The VSEPR model is widely used to predict molecular geometries and understand the behavior of molecules.

To apply the VSEPR model, we start by drawing the Lewis structure of the molecule. The Lewis structure shows the arrangement of atoms and valence electrons in a molecule. By counting the number of valence electrons and considering the octet rule , we can determine the Lewis structure of a molecule.

Application of VSEPR Model to NO2 Lewis Structure

Now let’s apply the VSEPR model to the NO2 molecule. NO2, also known as nitrogen dioxide, consists of one nitrogen atom (N) and two oxygen atoms (O).

To determine the Lewis structure of NO2, we first calculate the total number of valence electrons. Nitrogen has 5 valence electrons, and each oxygen atom has 6 valence electrons. Therefore, the total number of valence electrons in NO2 is:

5 (from nitrogen) + 2 * 6 (from oxygen) = 17

Next, we arrange the atom s in the molecule and connect them with single bonds. In the case of NO2, nitrogen is the central atom, and the two oxygen atoms are bonded to it.

To distribute the valence electrons, we place them around the atom s, starting with the outer atoms and then the central atom . In NO2, each oxygen atom needs 2 electrons to complete its octet, while nitrogen needs 3 electrons . This leaves us with 17 – 4 = 1 3 electrons to distribute.

We place the remaining electrons as l one pair s on the oxygen atoms. Each oxygen atom will have one l one pair , and nitrogen will have one l one pair as well.

Electron Geometry of NO2

The electron geometry of a molecule is determined by the arrangement of electron pairs around the central atom. In the case of NO2, nitrogen is the central atom, and it has one l one pair and two bonding pairs .

According to the VSEPR model, the presence of one l one pair and two bonding pairs gives NO2 an electron pair geometry of trigonal planar . This means that the electron pairs are arranged in a flat, triangular shape around the nitrogen atom.

The bond angle in NO2 is approximately 134 degrees. This angle is slightly less than the ideal bond angle of 120 degrees due to the repulsion between the l one pair and the bonding pairs.

In summary, the VSEPR model can be used to determine the electron geometry of NO2, which is trigonal planar . The presence of one l one pair and two bonding pairs results in a bond angle of approximately 134 degrees. Understanding the electron geometry of molecules like NO2 is crucial in predicting their physical and chemical properties .

Uses of NO2

Nitrogen dioxide (NO2) is a highly reactive and toxic gas that is commonly used in various industrial processes and applications. Its unique properties make it valuable for a range of purposes. Let’s explore some of the key uses of NO2.

Industrial production of Nitric acid

One of the primary uses of NO2 is in the industrial production of nitric acid. Nitric acid is a vital chemical compound used in the manufacturing of fertilizers, explosives, dyes, and pharmaceuticals. NO2 is a key intermediate in the Ostwald process , which involves the oxidation of ammonia to produce nitric acid. In this process , NO2 reacts with water to form nitric acid and nitrogen monoxide (NO). The production of nitric acid is crucial for various industries, making NO2 an essential component in its synthesis .

Catalyst in chemical reactions

NO2 also serves as a catalyst in several chemical reactions . A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process . NO2 acts as a catalyst in the oxidation of sulfur dioxide (SO2) to form sulfur trioxide (SO3). This reaction is a crucial step in the production of sulfuric acid, which is widely used in the manufacturing of fertilizers, detergents, and other chemical processes . The presence of NO2 enhances the efficiency of the reaction , making it an important catalyst in the sulfuric acid production process .

Regulation of sulfuric acid production

In addition to its role as a catalyst, NO2 also plays a significant role in the regulation of sulfuric acid production . The concentration of NO2 in the reaction mixture affects the rate of the oxidation reaction. By controlling the amount of NO2 present, manufacturers can regulate the production of sulfuric acid and ensure optimal efficiency . This regulation is crucial for maintaining the quality and quantity of sulfuric acid produced, as well as minimizing the environment al impact of the process .

Use as an oxidizer in rocket fuels

NO2 finds application as an oxidizer in rocket fuels . In rocket propulsion systems , an oxidizer is required to support the combustion of the fuel . NO2 is a powerful oxidizer that provides the necessary oxygen for the combustion process. It is commonly used in combination with fuels such as hydrazine to create highly energetic propulsion systems . The use of NO2 as an oxidizer allows rockets to achieve high speeds and propel payloads into space.

Manufacture of oxidized cellulose compounds

NO2 is also utilized in the manufacture of oxidized cellulose compounds . Oxidized cellulose is a material derived from cellulose fibers that have been chemically modified to enhance their properties . NO2 is used in the oxidation process, which introduces functional groups onto the cellulose structure , resulting in improved strength , stability, and absorbency. Oxidized cellulose compounds find applications in various industries, including healthcare, textiles, and paper manufacturing .

In conclusion, NO2 has a wide range of uses in various industries and applications. From its role in the production of nitric acid and regulation of sulfuric acid production to its use as a catalyst and oxidizer, NO2 plays a crucial role in numerous chemical processes . Its unique properties make it a valuable resource for enhancing the efficiency and performance of various industrial applications . Conclusion

In conclusion, understanding the NO2 Lewis structure is crucial for comprehending the chemical properties and behavior of nitrogen dioxide. By examining the arrangement of atoms and electrons within the molecule, we can gain insights into its polarity , reactivity, and overall stability . The Lewis structure of NO2 reveals that it consists of a central nitrogen atom bonded to two oxygen atoms, with one of the oxygen atoms carrying an unshared electron pair . This arrangement gives rise to a bent molecular geometry , resulting in a polar molecule with a partial positive charge on the nitrogen atom and partial negative charges on the oxygen atoms. The presence of a l one pair on one of the oxygen atoms makes NO2 highly reactive, particularly in terms of its involvement in atmospheric chemistry and air pollution . By studying the NO2 Lewis structure, scientists can better understand the behavior of this important compound and its impact on the environment and human health .

Frequently Asked Questions

1. how do you determine the lewis structure of no2.

To determine the Lewis structure of NO2 (nitrogen dioxide), you need to count the valence electrons and follow the octet rule . The Lewis structure for NO2 consists of a nitrogen atom bonded to two oxygen atoms, with a double bond between nitrogen and one oxygen atom, and a single bond between nitrogen and the other oxygen atom.

2. What is the hybridization of NO2?

The hybridization of NO2 (nitrogen dioxide) is sp2. In the Lewis structure of NO2, the nitrogen atom forms three sigm a bond s with the two oxygen atoms and has one l one pair . This arrangement requires the nitrogen atom to undergo sp2 hybridization.

3. How many valence electrons are in the Lewis structure of NO2-?

In the Lewis structure of NO2- ( nitrite ion ), there are 1 8 valence electrons . The nitrogen atom contributes 5 valence electrons, and each oxygen atom contributes 6 valence electrons. The negative charge on the nitrite ion adds an additional electron , totaling 1 8 valence electrons .

4. Does the Lewis structure of NO2 follow the octet rule?

Yes, the Lewis structure of NO2 (nitrogen dioxide) follows the octet rule . The nitrogen atom has a double bond with one oxygen atom and a single bond with the other oxygen atom, resulting in a total of 8 valence electrons around the nitrogen atom.

5. Why does NO2 have a double bond?

NO2 (nitrogen dioxide) has a double bond because it allows the nitrogen atom to achieve a stable octet configuration . By forming a double bond with one of the oxygen atoms, the nitrogen atom can share two pairs of electrons, satisfying the octet rule .

6. What is the bond order in the Lewis structure of NO2+?

The bond order in the Lewis structure of NO2+ (nitronium ion ) is 2. The nitrogen atom forms a double bond with one of the oxygen atoms and a coordinate covalent bond with the other oxygen atom, resulting in a bond order of 2.

7. How do you draw the Lewis structure of NO2?

To draw the Lewis structure of NO2 (nitrogen dioxide), start by placing the nitrogen atom in the center. Connect the nitrogen atom to two oxygen atoms using single bonds. Then, add a double bond between the nitrogen atom and one of the oxygen atoms. Finally, distribute any remaining valence electrons as l one pair s.

8. Does the Lewis structure of NO2 exhibit resonance?

Yes, the Lewis structure of NO2 ( nitrogen dioxide) exhibits resonance . The double bond in the structure can be delocalized between the nitrogen atom and either of the oxygen atoms, resulting in resonance structures .

9. What is the bond angle in the Lewis structure of NO2?

The bond angle in the Lewis structure of NO2 (nitrogen dioxide) is approximately 134 degrees. The oxygen atoms are arranged in a bent shape around the nitrogen atom, resulting in a bond angle slightly less than 180 degrees .

10. Is the Lewis structure of NO2 polar or nonpolar?

The Lewis structure of NO2 (nitrogen dioxide) is polar. The presence of a bent molecular geometry and the unequal distribution of electrons due to the double bond result in a polar molecule.

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sarnali mukherjee

Hi…..I am Sarnali Mukherjee, a graduate from the University of Calcutta. I love to teach and share knowledge on chemistry. I have gradually gained interest in article writing since one year ago. I would love to acquire more knowledge on my subject in the future. Let’s connect through LinkedIn:

7.4 Formal Charges and Resonance

Learning objectives.

By the end of this section, you will be able to:

  • Compute formal charges for atoms in any Lewis structure
  • Use formal charges to identify the most reasonable Lewis structure for a given molecule
  • Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule

In the previous section, we discussed how to write Lewis structures for molecules and polyatomic ions. As we have seen, however, in some cases, there is seemingly more than one valid structure for a molecule. We can use the concept of formal charges to help us predict the most appropriate Lewis structure when more than one is reasonable.

Calculating Formal Charge

The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.

Thus, we calculate formal charge as follows:

We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.

We must remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.

Example 7.6

Calculating formal charge from lewis structures.

  • Step 2. We assign lone pairs of electrons to their atoms . Each Cl atom now has seven electrons assigned to it, and the I atom has eight.
  • Step 3. Subtract this number from the number of valence electrons for the neutral atom: I: 7 – 8 = –1 Cl: 7 – 7 = 0 The sum of the formal charges of all the atoms equals –1, which is identical to the charge of the ion (–1).

Check Your Learning

Example 7.7.

  • Step 2. Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.
  • Step 3. Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge: Br: 7 – 7 = 0 Cl: 7 – 7 = 0 All atoms in BrCl 3 have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.

N: 0; all three Cl atoms: 0

Using Formal Charge to Predict Molecular Structure

The arrangement of atoms in a molecule or ion is called its molecular structure . In many cases, following the steps for writing Lewis structures may lead to more than one possible molecular structure—different multiple bond and lone-pair electron placements or different arrangements of atoms, for instance. A few guidelines involving formal charge can be helpful in deciding which of the possible structures is most likely for a particular molecule or ion:

  • A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.
  • If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.
  • Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.
  • When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

To see how these guidelines apply, let us consider some possible structures for carbon dioxide, CO 2 . We know from our previous discussion that the less electronegative atom typically occupies the central position, but formal charges allow us to understand why this occurs. We can draw three possibilities for the structure: carbon in the center and double bonds, carbon in the center with a single and triple bond, and oxygen in the center with double bonds:

Comparing the three formal charges, we can definitively identify the structure on the left as preferable because it has only formal charges of zero (Guideline 1).

As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular structures: NCS – , CNS – , or CSN – . The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:

Note that the sum of the formal charges in each case is equal to the charge of the ion (–1). However, the first arrangement of atoms is preferred because it has the lowest number of atoms with nonzero formal charges (Guideline 2). Also, it places the least electronegative atom in the center, and the negative charge on the more electronegative element (Guideline 4).

Example 7.8

Using formal charge to determine molecular structure.

The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:

The number of atoms with formal charges are minimized (Guideline 2), there is no formal charge with a magnitude greater than one (Guideline 2), the negative formal charge is on the more electronegative element (Guideline 4), and the less electronegative atom is in the center position.

Notice that the more likely structure for the nitrite anion in Example 7.8 may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O bonds:

If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. Experiments show, however, that both N–O bonds in NO 2 − NO 2 − have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for NO 2 − NO 2 − in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of resonance : if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in NO 2 − NO 2 − is the average of a double bond and a single bond. We call the individual Lewis structures resonance forms . The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms.

We should remember that a molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).

The carbonate anion, CO 3 2− , CO 3 2− , provides a second example of resonance:

One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C–O bonds are exactly the same.

Link to Learning

Use this online quiz to practice your skills in drawing resonance structures and estimating formal charges.

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Home / A Key Skill: How to Calculate Formal Charge

Bonding, Structure, and Resonance

By James Ashenhurst

  • A Key Skill: How to Calculate Formal Charge

Last updated: May 1st, 2024 |

How To Calculate Formal Charge

To calculate the formal charge of an atom, we start by:

  • evaluating the number of valence electrons ( VE ) the neutral atom has (e.g. 3 for boron, 4 for carbon, 5 for nitrogen, and so on).  (note: this is also equivalent to the effective nuclear charge Z eff , the number of protons that an electron in the valence orbital “sees” due to screening by inner-shell electrons.)
  • counting the number of  non-bonded valence electrons ( NBE ) on the atom. Each lone pair counts as  2 , and each unpaired electron counts as 1.
  • counting the number of  bonds ( B ) to the atom, or alternatively, counting the number of bonding electrons and dividing this by 2 .

The formal charge  FC is then calculated by subtracting NBE  and B  from VE .

FC = VE – ( NBE + B ) 

which is equivalent to

FC = VE – NBE – B

The calculation is pretty straightforward if all the information is given to you. However, for brevity’s sake, there are many times when lone pairs and C-H bonds are not explicitly drawn out .

So part of the trick for you will be to calculate the formal charge in situations where you have to take account of implicit  lone pairs and C-H bonds.

In the article below, we’ll address many of these situations. We’ll also warn you of the situations where the calculated formal charge of an atom is not necessarily a good clue as to its reactivity , which is extremely important going forward.

Table of Contents

  • Formal Charge
  • Simple Examples For First-Row Elements
  • Formal Charge Calculations When You Aren’t Given All The Details
  • Some Classic Formal Charge Problems
  • Formal Charges and Curved Arrows

Quiz Yourself!

(advanced) references and further reading, 1. formal charge.

Formal charge is a book-keeping formalism for assigning a charge to a specific atom.

To obtain the formal charge of an atom, we start by counting the number of valence electrons [ Note 1 ] for the neutral atom , and then subtract from it the number of electrons that it “ owns ” ( i.e. electrons in lone pairs, or singly-occupied orbitals ) and half of the electrons that it shares ( half the number of bonding electrons, which is equivalent to the number of bonds )

The simplest way to write the formula for formal charge  ( FC)  is:

FC = VE – NBE – B

  • VE corresponds to the number of electrons around the neutral atom (3 for boron, 4 for carbon, 5 for nitrogen, 6 for oxygen, 7 for fluorine)
  • NBE corresponds to the number of non-bonded electrons around the atom (2 for a lone pair, 1 for a singly-occupied orbital, 0 for an empty orbital)
  • B is the number of  bonds around the atom (equivalent to half the number of bonding electrons)

It’s called “ formal ” charge because it assumes that all bonding electrons are shared equally . It doesn’t account for electronegativity differences (i.e. dipoles).

For that reason formal charge isn’t always a good guide to where the electrons actually are in a molecule and can be an unreliable guide to reactivity. We’ll have more to say on that below .

2. Simple Examples For First-Row Elements

When all the lone pairs are drawn out for you, calculating formal charge is fairly straightforward.

Let’s work through the first example in the quiz below.

  • In the hydronium ion (H 3 O) the central atom is oxygen , which has 6 valence electrons in the neutral atom
  • The central atom has 2 unpaired electrons and 3 bonds
  • The formal charge on oxygen is [6 – 2 – 3 = +1 ] giving us H 3 O +

See if you can fill in the rest for the examples below.

If that went well, you could try filling in the formal charges for all of the examples in this table.

Become a member to see the clickable quiz with answers on the back.

It will take some getting used to formal charge, but after a period of time it will be  assumed that you understand how to calculate formal charge, and that you can recognize structures where atoms will have a formal charge.

Let’s deal with some slightly trickier cases.

3. Formal Charge Calculations When You Aren’t Given All The Details

When we draw a stick figure of a person and don’t draw in their fingers, it doesn’t mean we’re drawing someone who had a bad day working with a table saw . We just assume that you could fill in the fingers if you really needed to, but you’re skipping it just to save time.

Chemical line drawings are like stick figures. They omit a lot of detail but still assume you know that certain things are there.

  • With carbon, we often omit drawing hydrogens . You’re still supposed to know that they are there, and add as many hydrogens as necessary to give a full octet (or sextet, if it’s a carbocation). 
  • If there is a lone pair or unpaired electron on a carbon, it’s always drawn in .

One note. If we draw a stick figure, and we do draw the fingers, and took the time to only draw in only 3 , then we can safely assume that the person really does only have 3 fingers . So in  the last two examples on that quiz we had to draw in the hydrogens in order for you to know that it was a carbocation, otherwise you would have to assume that it had a full octet!  

Oxygen and nitrogen (and the halogens) are dealt with slightly differently.

  • Bonds to hydrogen are always drawn in.
  • The lone pairs that are often omitted.
  • Nitrogen and oxygen will always have full octets. Always. [ Note 2 – OK, two exceptions ]

So even when the lone pairs aren’t drawn in, assume that enough are present to make a full octet . And when bonds from these atoms to hydrogen are missing , that means exactly what it seems to be: there really isn’t any hydrogen!

Try these examples:

Now see if you can put these examples together!

(Note that some of these are not stable molecules, but instead represent are resonance forms that you will encounter at various points during the course!)

4. Some Classic Formal Charge Questions

We can use the exact same formal charge formula, above, along with the rules for implicit lone pairs and hydrogens, to figure out the formal charge of atoms in some pretty exotic-looking molecules.

Here are some classic formal charge problems.

The formal charge formula can even be applied to some fairly exotic reactive intermediates we’ll meet later in the semester.

Don’t get spooked out. Just count the electrons and the bonds, and that will lead you to the right answer.

5. Formal Charges and Curved Arrows

We use curved arrows to show the movement of electron pairs in reactions and in resonance structures. ( See post: Curved Arrows For Reactions )

For example, here is a curved arrow that shows the reaction of the hydroxide ion HO(-) with a proton (H+).

The arrow shows movement of two electrons from oxygen to form a new O–H bond .

Curved arrows are also useful for keeping track of changes in formal charge.  Note that the formal charge at the initial tail of the curved arrow (the oxygen) becomes more positive (from -1 to 0) and the formal charge at the final tail (the H+) becomes more negative (from +1 to 0). 

When acid is added to water, we form the hydronium ion, H 3 O + .

Here’s a quiz. See if you can draw the curved arrow going from the hydroxide ion to H 3 O+.

If you did it successfully – congratulations!

But I’m willing to bet that at least a small percentage of you drew the arrow going to the positively charged oxygen .

What’s wrong with that?

There isn’t an empty orbital on oxygen that can accept the lone pair.  If you follow the logic of curved arrows, that would result in a new O–O bond, and 10 electrons on the oxygen, breaking the octet rule.

Hold on a minute, you might say. “ I thought oxygen was positively charged? I f it doesn’t react on oxygen, where is it supposed to react ?”

On the hydrogens! H 3 O+ is Brønsted acid, after all. Right?

This is a great illustration of the reason why it’s called “ formal charge”, and how formal charge not the same as  electrostatic charge (a.ka. “partial charges” or “electron density”).

Formal charge is ultimately a book-keeping formalism, a little bit like assigning the “win” to one of the 5 pitchers in a baseball game. [ Note 3 ] It doesn’t take into account the fact that the electrons in the oxygen-hydrogen bond are unequally shared, with a substantial dipole.

So although we draw a “formal” charge on oxygen, the partial positive charges are all on  hydrogen. Despite bearing a positive formal charge bears a partially negative electrostatic charge.

This is why bases such as HO(-) react at the H, not the oxygen.

Just to reiterate:

  • Positive charges on oxygen and nitrogen do not represent an empty orbital. Assume that oxygen and nitrogen have full octets! [ Note 2 ]
  • In contrast, positive charges on carbon do represent empty orbitals.

6. Halogens

Positive formal charges on halogens fall into two main categories.

We’ll often be found drawing  halonium ions   Cl+ , Br+, and I+ as species with six valence electrons and an empty orbital  ( but never F+ – it’s a ravenous beast )

It’s OK to think of these species as bearing an empty orbital since they are large and relatively polarizable .  They can distribute the positive charge over their relatively large volume.

These species can accept a lone pair of electrons from a Lewis base, resulting in a full octet.

Cl, Br, and I can also bear positive formal charges as a result of being bonded to two atoms.

It’s important to realize in these cases that the halogen bears a  full octet and not an empty orbital. They will therefore not directly accept a pair of electrons from Lewis bases; it’s often the case that the atom adjacent to the halogen accepts the electrons.

7. Conclusion

If you have reached the end and did all the quizzes, you should be well prepared for all the examples of formal charge you see in the rest of the course.

  • Formal charge can be calculated using the formula FC = VE – NBE – B
  • Line drawings often omit lone pairs and C-H bonds. Be alert for these situations when calculating formal charges.
  • Positively charged carbon has an empty orbital, but assume that positively charged nitrogen and oxygen have full octets.
  • The example of the hydronium ion H 3 O+ shows the perils of relying on formal charge to understand reactivity. Pay close attention to the differences in electronegativity between atoms and draw out the dipoles to get a true sense of their reactivity.

Related Articles

  • Partial Charges Give Clues About Electron Flow
  • How To Use Electronegativity To Determine Electron Density (and why NOT to trust formal charge)
  • How to apply electronegativity and resonance to understand reactivity
  • Maybe they should call them, “Formal Wins” ?
  • Common Mistakes: Formal Charges Can Mislead

Note 1. Using “valence electrons” gets you the right answer. But if you think about it, it doesn’t quite make sense. Where do positive charges come from? From the positively charged protons in the nucleus, of course!

So the “valence electrons” part of this equation is more properly thought of as a proxy for valence protons – which is another way of saying the “ effective nuclear charge” ; the charge felt by each valence electron from the nucleus, not counting the filled inner shells.

Note 2. Nitrenes are an exception. Another exception is when we want to draw  bad resonance forms.

Note 3 . In baseball, every game results in a win or a loss for the team . Back in the days of   Old Hoss Radborn , where complete games were the norm, a logical extension of this was to assign the win to the individual pitcher. In today’s era, with multiple relief pitchers, there are rules for determining which pitcher gets credited with the win. It’s very possible for a pitcher to get completely shelled on the mound and yet, through fortuitous circumstance, still be credited for the win.  See post: Maybe They Should Call Them, “Formal Wins” ? 

In the same way, oxygen is given individual credit for the charge of +1 on the hydronium ion , H 3 O+, even though the actual positive electrostatic charge is distributed among the hydrogens.

Note 4. This image from a previous incarnation of this post demonstates some relationships for the geometry of various compounds of first-row elements.

1. Valence, Oxidation Number, and Formal Charge: Three Related but Fundamentally Different Concepts Gerard Parkin Journal of Chemical Education 2006 83 (5), 791 DOI : 10.1021/ed083p791 

2. Lewis structures, formal charge, and oxidation numbers: A more user-friendly approach John E. Packer and Sheila D. Woodgate Journal of Chemical Education   1991   68  (6), 456 DOI : 10.1021/ed068p456

00 General Chemistry Review

  • Lewis Structures
  • Ionic and Covalent Bonding
  • Chemical Kinetics
  • Chemical Equilibria
  • Valence Electrons of the First Row Elements
  • How Concepts Build Up In Org 1 ("The Pyramid")

01 Bonding, Structure, and Resonance

  • How Do We Know Methane (CH4) Is Tetrahedral?
  • Hybrid Orbitals and Hybridization
  • How To Determine Hybridization: A Shortcut
  • Orbital Hybridization And Bond Strengths
  • Sigma bonds come in six varieties: Pi bonds come in one
  • The Four Intermolecular Forces and How They Affect Boiling Points
  • 3 Trends That Affect Boiling Points
  • Introduction to Resonance
  • How To Use Curved Arrows To Interchange Resonance Forms
  • Evaluating Resonance Forms (1) - The Rule of Least Charges
  • How To Find The Best Resonance Structure By Applying Electronegativity
  • Evaluating Resonance Structures With Negative Charges
  • Evaluating Resonance Structures With Positive Charge
  • Exploring Resonance: Pi-Donation
  • Exploring Resonance: Pi-acceptors
  • In Summary: Evaluating Resonance Structures
  • Drawing Resonance Structures: 3 Common Mistakes To Avoid
  • Bond Hybridization Practice
  • Structure and Bonding Practice Quizzes
  • Resonance Structures Practice

02 Acid Base Reactions

  • Introduction to Acid-Base Reactions
  • Acid Base Reactions In Organic Chemistry
  • The Stronger The Acid, The Weaker The Conjugate Base
  • Walkthrough of Acid-Base Reactions (3) - Acidity Trends
  • Five Key Factors That Influence Acidity
  • Acid-Base Reactions: Introducing Ka and pKa
  • How to Use a pKa Table
  • The pKa Table Is Your Friend
  • A Handy Rule of Thumb for Acid-Base Reactions
  • Acid Base Reactions Are Fast
  • pKa Values Span 60 Orders Of Magnitude
  • How Protonation and Deprotonation Affect Reactivity
  • Acid Base Practice Problems

03 Alkanes and Nomenclature

  • Meet the (Most Important) Functional Groups
  • Condensed Formulas: Deciphering What the Brackets Mean
  • Hidden Hydrogens, Hidden Lone Pairs, Hidden Counterions
  • Don't Be Futyl, Learn The Butyls
  • Primary, Secondary, Tertiary, Quaternary In Organic Chemistry
  • Branching, and Its Affect On Melting and Boiling Points
  • The Many, Many Ways of Drawing Butane
  • Wedge And Dash Convention For Tetrahedral Carbon
  • Common Mistakes in Organic Chemistry: Pentavalent Carbon
  • Table of Functional Group Priorities for Nomenclature
  • Summary Sheet - Alkane Nomenclature
  • Organic Chemistry IUPAC Nomenclature Demystified With A Simple Puzzle Piece Approach
  • Boiling Point Quizzes
  • Organic Chemistry Nomenclature Quizzes

04 Conformations and Cycloalkanes

  • Staggered vs Eclipsed Conformations of Ethane
  • Conformational Isomers of Propane
  • Newman Projection of Butane (and Gauche Conformation)
  • Introduction to Cycloalkanes (1)
  • Geometric Isomers In Small Rings: Cis And Trans Cycloalkanes
  • Calculation of Ring Strain In Cycloalkanes
  • Cycloalkanes - Ring Strain In Cyclopropane And Cyclobutane
  • Cyclohexane Conformations
  • Cyclohexane Chair Conformation: An Aerial Tour
  • How To Draw The Cyclohexane Chair Conformation
  • The Cyclohexane Chair Flip
  • The Cyclohexane Chair Flip - Energy Diagram
  • Substituted Cyclohexanes - Axial vs Equatorial
  • Ranking The Bulkiness Of Substituents On Cyclohexanes: "A-Values"
  • Cyclohexane Chair Conformation Stability: Which One Is Lower Energy?
  • Fused Rings - Cis-Decalin and Trans-Decalin
  • Naming Bicyclic Compounds - Fused, Bridged, and Spiro
  • Bredt's Rule (And Summary of Cycloalkanes)
  • Newman Projection Practice
  • Cycloalkanes Practice Problems

05 A Primer On Organic Reactions

  • The Most Important Question To Ask When Learning a New Reaction
  • Learning New Reactions: How Do The Electrons Move?
  • The Third Most Important Question to Ask When Learning A New Reaction
  • 7 Factors that stabilize negative charge in organic chemistry
  • 7 Factors That Stabilize Positive Charge in Organic Chemistry
  • Nucleophiles and Electrophiles
  • Curved Arrows (for reactions)
  • Curved Arrows (2): Initial Tails and Final Heads
  • Nucleophilicity vs. Basicity
  • The Three Classes of Nucleophiles
  • What Makes A Good Nucleophile?
  • What makes a good leaving group?
  • 3 Factors That Stabilize Carbocations
  • Equilibrium and Energy Relationships
  • What's a Transition State?
  • Hammond's Postulate
  • Learning Organic Chemistry Reactions: A Checklist (PDF)
  • Introduction to Free Radical Substitution Reactions
  • Introduction to Oxidative Cleavage Reactions

06 Free Radical Reactions

  • Bond Dissociation Energies = Homolytic Cleavage
  • Free Radical Reactions
  • 3 Factors That Stabilize Free Radicals
  • What Factors Destabilize Free Radicals?
  • Bond Strengths And Radical Stability
  • Free Radical Initiation: Why Is "Light" Or "Heat" Required?
  • Initiation, Propagation, Termination
  • Monochlorination Products Of Propane, Pentane, And Other Alkanes
  • Selectivity In Free Radical Reactions
  • Selectivity in Free Radical Reactions: Bromination vs. Chlorination
  • Halogenation At Tiffany's
  • Allylic Bromination
  • Bonus Topic: Allylic Rearrangements
  • In Summary: Free Radicals
  • Synthesis (2) - Reactions of Alkanes
  • Free Radicals Practice Quizzes

07 Stereochemistry and Chirality

  • Types of Isomers: Constitutional Isomers, Stereoisomers, Enantiomers, and Diastereomers
  • How To Draw The Enantiomer Of A Chiral Molecule
  • How To Draw A Bond Rotation
  • Introduction to Assigning (R) and (S): The Cahn-Ingold-Prelog Rules
  • Assigning Cahn-Ingold-Prelog (CIP) Priorities (2) - The Method of Dots
  • Enantiomers vs Diastereomers vs The Same? Two Methods For Solving Problems
  • Assigning R/S To Newman Projections (And Converting Newman To Line Diagrams)
  • How To Determine R and S Configurations On A Fischer Projection
  • The Meso Trap
  • Optical Rotation, Optical Activity, and Specific Rotation
  • Optical Purity and Enantiomeric Excess
  • What's a Racemic Mixture?
  • Chiral Allenes And Chiral Axes
  • Stereochemistry Practice Problems and Quizzes

08 Substitution Reactions

  • Introduction to Nucleophilic Substitution Reactions
  • Walkthrough of Substitution Reactions (1) - Introduction
  • Two Types of Nucleophilic Substitution Reactions
  • The SN2 Mechanism
  • Why the SN2 Reaction Is Powerful
  • The SN1 Mechanism
  • The Conjugate Acid Is A Better Leaving Group
  • Comparing the SN1 and SN2 Reactions
  • Polar Protic? Polar Aprotic? Nonpolar? All About Solvents
  • Steric Hindrance is Like a Fat Goalie
  • Common Blind Spot: Intramolecular Reactions
  • The Conjugate Base is Always a Stronger Nucleophile
  • Substitution Practice - SN1
  • Substitution Practice - SN2

09 Elimination Reactions

  • Elimination Reactions (1): Introduction And The Key Pattern
  • Elimination Reactions (2): The Zaitsev Rule
  • Elimination Reactions Are Favored By Heat
  • Two Elimination Reaction Patterns
  • The E1 Reaction
  • The E2 Mechanism
  • E1 vs E2: Comparing the E1 and E2 Reactions
  • Antiperiplanar Relationships: The E2 Reaction and Cyclohexane Rings
  • Bulky Bases in Elimination Reactions
  • Comparing the E1 vs SN1 Reactions
  • Elimination (E1) Reactions With Rearrangements
  • E1cB - Elimination (Unimolecular) Conjugate Base
  • Elimination (E1) Practice Problems And Solutions
  • Elimination (E2) Practice Problems and Solutions

10 Rearrangements

  • Introduction to Rearrangement Reactions
  • Rearrangement Reactions (1) - Hydride Shifts
  • Carbocation Rearrangement Reactions (2) - Alkyl Shifts
  • Pinacol Rearrangement
  • The SN1, E1, and Alkene Addition Reactions All Pass Through A Carbocation Intermediate

11 SN1/SN2/E1/E2 Decision

  • Identifying Where Substitution and Elimination Reactions Happen
  • Deciding SN1/SN2/E1/E2 (1) - The Substrate
  • Deciding SN1/SN2/E1/E2 (2) - The Nucleophile/Base
  • SN1 vs E1 and SN2 vs E2 : The Temperature
  • Deciding SN1/SN2/E1/E2 - The Solvent
  • Wrapup: The Key Factors For Determining SN1/SN2/E1/E2
  • Alkyl Halide Reaction Map And Summary
  • SN1 SN2 E1 E2 Practice Problems

12 Alkene Reactions

  • E and Z Notation For Alkenes (+ Cis/Trans)
  • Alkene Stability
  • Alkene Addition Reactions: "Regioselectivity" and "Stereoselectivity" (Syn/Anti)
  • Stereoselective and Stereospecific Reactions
  • Hydrohalogenation of Alkenes and Markovnikov's Rule
  • Hydration of Alkenes With Aqueous Acid
  • Rearrangements in Alkene Addition Reactions
  • Halogenation of Alkenes and Halohydrin Formation
  • Oxymercuration Demercuration of Alkenes
  • Hydroboration Oxidation of Alkenes
  • m-CPBA (meta-chloroperoxybenzoic acid)
  • OsO4 (Osmium Tetroxide) for Dihydroxylation of Alkenes
  • Palladium on Carbon (Pd/C) for Catalytic Hydrogenation of Alkenes
  • Cyclopropanation of Alkenes
  • A Fourth Alkene Addition Pattern - Free Radical Addition
  • Alkene Reactions: Ozonolysis
  • Summary: Three Key Families Of Alkene Reaction Mechanisms
  • Synthesis (4) - Alkene Reaction Map, Including Alkyl Halide Reactions
  • Alkene Reactions Practice Problems

13 Alkyne Reactions

  • Acetylides from Alkynes, And Substitution Reactions of Acetylides
  • Partial Reduction of Alkynes With Lindlar's Catalyst
  • Partial Reduction of Alkynes With Na/NH3 To Obtain Trans Alkenes
  • Alkyne Hydroboration With "R2BH"
  • Hydration and Oxymercuration of Alkynes
  • Hydrohalogenation of Alkynes
  • Alkyne Halogenation: Bromination, Chlorination, and Iodination of Alkynes
  • Alkyne Reactions - The "Concerted" Pathway
  • Alkenes To Alkynes Via Halogenation And Elimination Reactions
  • Alkynes Are A Blank Canvas
  • Synthesis (5) - Reactions of Alkynes
  • Alkyne Reactions Practice Problems With Answers

14 Alcohols, Epoxides and Ethers

  • Alcohols - Nomenclature and Properties
  • Alcohols Can Act As Acids Or Bases (And Why It Matters)
  • Alcohols - Acidity and Basicity
  • The Williamson Ether Synthesis
  • Ethers From Alkenes, Tertiary Alkyl Halides and Alkoxymercuration
  • Alcohols To Ethers via Acid Catalysis
  • Cleavage Of Ethers With Acid
  • Epoxides - The Outlier Of The Ether Family
  • Opening of Epoxides With Acid
  • Epoxide Ring Opening With Base
  • Making Alkyl Halides From Alcohols
  • Tosylates And Mesylates
  • PBr3 and SOCl2
  • Elimination Reactions of Alcohols
  • Elimination of Alcohols To Alkenes With POCl3
  • Alcohol Oxidation: "Strong" and "Weak" Oxidants
  • Demystifying The Mechanisms of Alcohol Oxidations
  • Protecting Groups For Alcohols
  • Thiols And Thioethers
  • Calculating the oxidation state of a carbon
  • Oxidation and Reduction in Organic Chemistry
  • Oxidation Ladders
  • SOCl2 Mechanism For Alcohols To Alkyl Halides: SN2 versus SNi
  • Alcohol Reactions Roadmap (PDF)
  • Alcohol Reaction Practice Problems
  • Epoxide Reaction Quizzes
  • Oxidation and Reduction Practice Quizzes

15 Organometallics

  • What's An Organometallic?
  • Formation of Grignard and Organolithium Reagents
  • Organometallics Are Strong Bases
  • Reactions of Grignard Reagents
  • Protecting Groups In Grignard Reactions
  • Synthesis Problems Involving Grignard Reagents
  • Grignard Reactions And Synthesis (2)
  • Organocuprates (Gilman Reagents): How They're Made
  • Gilman Reagents (Organocuprates): What They're Used For
  • The Heck, Suzuki, and Olefin Metathesis Reactions (And Why They Don't Belong In Most Introductory Organic Chemistry Courses)
  • Reaction Map: Reactions of Organometallics
  • Grignard Practice Problems

16 Spectroscopy

  • Degrees of Unsaturation (or IHD, Index of Hydrogen Deficiency)
  • Conjugation And Color (+ How Bleach Works)
  • Introduction To UV-Vis Spectroscopy
  • UV-Vis Spectroscopy: Absorbance of Carbonyls
  • UV-Vis Spectroscopy: Practice Questions
  • Bond Vibrations, Infrared Spectroscopy, and the "Ball and Spring" Model
  • Infrared Spectroscopy: A Quick Primer On Interpreting Spectra
  • IR Spectroscopy: 4 Practice Problems
  • 1H NMR: How Many Signals?
  • Homotopic, Enantiotopic, Diastereotopic
  • Diastereotopic Protons in 1H NMR Spectroscopy: Examples
  • C13 NMR - How Many Signals
  • Liquid Gold: Pheromones In Doe Urine
  • Natural Product Isolation (1) - Extraction
  • Natural Product Isolation (2) - Purification Techniques, An Overview
  • Structure Determination Case Study: Deer Tarsal Gland Pheromone

17 Dienes and MO Theory

  • What To Expect In Organic Chemistry 2
  • Are these molecules conjugated?
  • Conjugation And Resonance In Organic Chemistry
  • Bonding And Antibonding Pi Orbitals
  • Molecular Orbitals of The Allyl Cation, Allyl Radical, and Allyl Anion
  • Pi Molecular Orbitals of Butadiene
  • Reactions of Dienes: 1,2 and 1,4 Addition
  • Thermodynamic and Kinetic Products
  • More On 1,2 and 1,4 Additions To Dienes
  • s-cis and s-trans
  • The Diels-Alder Reaction
  • Cyclic Dienes and Dienophiles in the Diels-Alder Reaction
  • Stereochemistry of the Diels-Alder Reaction
  • Exo vs Endo Products In The Diels Alder: How To Tell Them Apart
  • HOMO and LUMO In the Diels Alder Reaction
  • Why Are Endo vs Exo Products Favored in the Diels-Alder Reaction?
  • Diels-Alder Reaction: Kinetic and Thermodynamic Control
  • The Retro Diels-Alder Reaction
  • The Intramolecular Diels Alder Reaction
  • Regiochemistry In The Diels-Alder Reaction
  • The Cope and Claisen Rearrangements
  • Electrocyclic Reactions
  • Electrocyclic Ring Opening And Closure (2) - Six (or Eight) Pi Electrons
  • Diels Alder Practice Problems
  • Molecular Orbital Theory Practice

18 Aromaticity

  • Introduction To Aromaticity
  • Rules For Aromaticity
  • Huckel's Rule: What Does 4n+2 Mean?
  • Aromatic, Non-Aromatic, or Antiaromatic? Some Practice Problems
  • Antiaromatic Compounds and Antiaromaticity
  • The Pi Molecular Orbitals of Benzene
  • The Pi Molecular Orbitals of Cyclobutadiene
  • Frost Circles
  • Aromaticity Practice Quizzes

19 Reactions of Aromatic Molecules

  • Electrophilic Aromatic Substitution: Introduction
  • Activating and Deactivating Groups In Electrophilic Aromatic Substitution
  • Electrophilic Aromatic Substitution - The Mechanism
  • Ortho-, Para- and Meta- Directors in Electrophilic Aromatic Substitution
  • Understanding Ortho, Para, and Meta Directors
  • Why are halogens ortho- para- directors?
  • Disubstituted Benzenes: The Strongest Electron-Donor "Wins"
  • Electrophilic Aromatic Substitutions (1) - Halogenation of Benzene
  • Electrophilic Aromatic Substitutions (2) - Nitration and Sulfonation
  • EAS Reactions (3) - Friedel-Crafts Acylation and Friedel-Crafts Alkylation
  • Intramolecular Friedel-Crafts Reactions
  • Nucleophilic Aromatic Substitution (NAS)
  • Nucleophilic Aromatic Substitution (2) - The Benzyne Mechanism
  • Reactions on the "Benzylic" Carbon: Bromination And Oxidation
  • The Wolff-Kishner, Clemmensen, And Other Carbonyl Reductions
  • More Reactions on the Aromatic Sidechain: Reduction of Nitro Groups and the Baeyer Villiger
  • Aromatic Synthesis (1) - "Order Of Operations"
  • Synthesis of Benzene Derivatives (2) - Polarity Reversal
  • Aromatic Synthesis (3) - Sulfonyl Blocking Groups
  • Birch Reduction
  • Synthesis (7): Reaction Map of Benzene and Related Aromatic Compounds
  • Aromatic Reactions and Synthesis Practice
  • Electrophilic Aromatic Substitution Practice Problems

20 Aldehydes and Ketones

  • What's The Alpha Carbon In Carbonyl Compounds?
  • Nucleophilic Addition To Carbonyls
  • Aldehydes and Ketones: 14 Reactions With The Same Mechanism
  • Sodium Borohydride (NaBH4) Reduction of Aldehydes and Ketones
  • Grignard Reagents For Addition To Aldehydes and Ketones
  • Wittig Reaction
  • Hydrates, Hemiacetals, and Acetals
  • Imines - Properties, Formation, Reactions, and Mechanisms
  • All About Enamines
  • Breaking Down Carbonyl Reaction Mechanisms: Reactions of Anionic Nucleophiles (Part 2)
  • Aldehydes Ketones Reaction Practice

21 Carboxylic Acid Derivatives

  • Nucleophilic Acyl Substitution (With Negatively Charged Nucleophiles)
  • Addition-Elimination Mechanisms With Neutral Nucleophiles (Including Acid Catalysis)
  • Basic Hydrolysis of Esters - Saponification
  • Transesterification
  • Proton Transfer
  • Fischer Esterification - Carboxylic Acid to Ester Under Acidic Conditions
  • Lithium Aluminum Hydride (LiAlH4) For Reduction of Carboxylic Acid Derivatives
  • LiAlH[Ot-Bu]3 For The Reduction of Acid Halides To Aldehydes
  • Di-isobutyl Aluminum Hydride (DIBAL) For The Partial Reduction of Esters and Nitriles
  • Amide Hydrolysis
  • Thionyl Chloride (SOCl2)
  • Diazomethane (CH2N2)
  • Carbonyl Chemistry: Learn Six Mechanisms For the Price Of One
  • Making Music With Mechanisms (PADPED)
  • Carboxylic Acid Derivatives Practice Questions

22 Enols and Enolates

  • Keto-Enol Tautomerism
  • Enolates - Formation, Stability, and Simple Reactions
  • Kinetic Versus Thermodynamic Enolates
  • Aldol Addition and Condensation Reactions
  • Reactions of Enols - Acid-Catalyzed Aldol, Halogenation, and Mannich Reactions
  • Claisen Condensation and Dieckmann Condensation
  • Decarboxylation
  • The Malonic Ester and Acetoacetic Ester Synthesis
  • The Michael Addition Reaction and Conjugate Addition
  • The Robinson Annulation
  • Haloform Reaction
  • The Hell–Volhard–Zelinsky Reaction
  • Enols and Enolates Practice Quizzes
  • The Amide Functional Group: Properties, Synthesis, and Nomenclature
  • Basicity of Amines And pKaH
  • 5 Key Basicity Trends of Amines
  • The Mesomeric Effect And Aromatic Amines
  • Nucleophilicity of Amines
  • Alkylation of Amines (Sucks!)
  • Reductive Amination
  • The Gabriel Synthesis
  • Some Reactions of Azides
  • The Hofmann Elimination
  • The Hofmann and Curtius Rearrangements
  • The Cope Elimination
  • Protecting Groups for Amines - Carbamates
  • The Strecker Synthesis of Amino Acids
  • Introduction to Peptide Synthesis
  • Reactions of Diazonium Salts: Sandmeyer and Related Reactions
  • Amine Practice Questions

24 Carbohydrates

  • D and L Notation For Sugars
  • Pyranoses and Furanoses: Ring-Chain Tautomerism In Sugars
  • What is Mutarotation?
  • Reducing Sugars
  • The Big Damn Post Of Carbohydrate-Related Chemistry Definitions
  • The Haworth Projection
  • Converting a Fischer Projection To A Haworth (And Vice Versa)
  • Reactions of Sugars: Glycosylation and Protection
  • The Ruff Degradation and Kiliani-Fischer Synthesis
  • Isoelectric Points of Amino Acids (and How To Calculate Them)
  • Carbohydrates Practice
  • Amino Acid Quizzes

25 Fun and Miscellaneous

  • A Gallery of Some Interesting Molecules From Nature
  • Screw Organic Chemistry, I'm Just Going To Write About Cats
  • On Cats, Part 1: Conformations and Configurations
  • On Cats, Part 2: Cat Line Diagrams
  • On Cats, Part 4: Enantiocats
  • On Cats, Part 6: Stereocenters
  • Organic Chemistry Is Shit
  • The Organic Chemistry Behind "The Pill"
  • Maybe they should call them, "Formal Wins" ?
  • Why Do Organic Chemists Use Kilocalories?
  • The Principle of Least Effort
  • Organic Chemistry GIFS - Resonance Forms
  • Reproducibility In Organic Chemistry
  • What Holds The Nucleus Together?
  • How Reactions Are Like Music
  • Organic Chemistry and the New MCAT

26 Organic Chemistry Tips and Tricks

  • Draw The Ugly Version First
  • Organic Chemistry Study Tips: Learn the Trends
  • The 8 Types of Arrows In Organic Chemistry, Explained
  • Top 10 Skills To Master Before An Organic Chemistry 2 Final
  • Common Mistakes with Carbonyls: Carboxylic Acids... Are Acids!
  • Planning Organic Synthesis With "Reaction Maps"
  • Alkene Addition Pattern #1: The "Carbocation Pathway"
  • Alkene Addition Pattern #2: The "Three-Membered Ring" Pathway
  • Alkene Addition Pattern #3: The "Concerted" Pathway
  • Number Your Carbons!
  • The 4 Major Classes of Reactions in Org 1
  • How (and why) electrons flow
  • Grossman's Rule
  • Three Exam Tips
  • A 3-Step Method For Thinking Through Synthesis Problems
  • Putting It Together
  • Putting Diels-Alder Products in Perspective
  • The Ups and Downs of Cyclohexanes
  • The Most Annoying Exceptions in Org 1 (Part 1)
  • The Most Annoying Exceptions in Org 1 (Part 2)
  • The Marriage May Be Bad, But the Divorce Still Costs Money
  • 9 Nomenclature Conventions To Know
  • Nucleophile attacks Electrophile

27 Case Studies of Successful O-Chem Students

  • Success Stories: How Corina Got The The "Hard" Professor - And Got An A+ Anyway
  • How Helena Aced Organic Chemistry
  • From a "Drop" To B+ in Org 2 – How A Hard Working Student Turned It Around
  • How Serge Aced Organic Chemistry
  • Success Stories: How Zach Aced Organic Chemistry 1
  • Success Stories: How Kari Went From C– to B+
  • How Esther Bounced Back From a "C" To Get A's In Organic Chemistry 1 And 2
  • How Tyrell Got The Highest Grade In Her Organic Chemistry Course
  • This Is Why Students Use Flashcards
  • Success Stories: How Stu Aced Organic Chemistry
  • How John Pulled Up His Organic Chemistry Exam Grades
  • Success Stories: How Nathan Aced Organic Chemistry (Without It Taking Over His Life)
  • How Chris Aced Org 1 and Org 2
  • Interview: How Jay Got an A+ In Organic Chemistry
  • How to Do Well in Organic Chemistry: One Student's Advice
  • "America's Top TA" Shares His Secrets For Teaching O-Chem
  • "Organic Chemistry Is Like..." - A Few Metaphors
  • How To Do Well In Organic Chemistry: Advice From A Tutor
  • Guest post: "I went from being afraid of tests to actually looking forward to them".

Comment section

60 thoughts on “ a key skill: how to calculate formal charge ”.

Hello, thanks for your wonderful posts on organic chemistry. It reallys helps me to recap org chem and I really like how you explain all these topics with a bit of humor.

That said, I think in this posts may be some typos: I think there are two typos in the solution of the last quiz of chapter 3 (ID 2310): (a) In the third task [C3H7N] of the quiz, there is just one electron on the negative charged carbon. Shouldn’t there be two electrons? (b) And in the fourth task [O-CH2] the sign of the formal charge of the carbon atom should be +1 (in the calculation). (c) Note 4 says “(…) of various compounds of first-row elements.” Aren’t the shown elements in the picture from the second row of the periodic table?

Thank you very much!

  • Pingback: A Key Skill: How to Calculate Formal Charge | Straight A Mindset

Your explanations and examples were clear and easy to understand. I appreciate the detailed step-by-step instructions, which made it easy to follow along and understand the concept. Thank you for taking the time to create this helpful resource

I think for Quiz ID: 2310, the formal charge for the carbon in the fourth molecule should be +1 instead of -1.

Fixed. Thanks for the spot!

Thank you so much sir. Finally i understood how to calculate the formal charge

Nice simple explanation

Great teaching , can I know where did u studied ??

Hi I am extremely confused. The two formulas for calculating FC that you provided are not the same and don’t produce the same results when I tried them out.

Formal charge = [# of valence electrons] – [electrons in lone pairs + 1/2 the number of bonding electrons]

Formal Charge = [# of valence electrons on atom] – [non-bonded electrons + number of bonds].

They do not produce the same result… If I have the formula BH4, and use the first formula provided to find FC of B, I would get:

(3) – (0 + 2) = +1

Using the second formula provided:

(3) – (0+4) = -1

Aren’t these formulas supposed to produce the same results? I am quite confused and I don’t know if I missed something.

Ah. I should have been more clear. The number of bonding electrons in BH4 equals 8, since each bond has two electrons and there are 4 B-H bonds. Half of this number equals 4. This should give you the same answer. I have updated the post to make this more explicit.

  • Pingback: Como posso calcular a cobrança formal? – CorujaSabia

That was the best i have seen but i have a problem with the formula,i think the side where the shared pair electrons came was suppose to be negative but then yours was positive,so am finfding it difficult to understand because the slides we were given by our lecturer shows that it was subtracted not added. i would love it when u explain it to me.

  • Pingback: ¿Cómo puedo calcular el cargo formal? – ElbuhoSabio

It was a very great explanation! Now I have a good concept about how to find formula charge. And also i am just a grade nine student so i want to say thank you for this.

  • Pingback: Come posso calcolare l’addebito formale? – GufoSaggio

YOU ARE THE BEST. I GOT THE HIGHEST MARK IN MY FIRST QUIZ, AND I KNOW THAT THROUGH THIS I WILL GET THE BEST IN MIDTERM AND FINAL. I want you guys to go on youtube and follow the steps. THANK YOU VERY MUCH.

I remember learning that in the cyanide ion, the carbon is nucleophilic because the formal negative charge is on carbon, not nitrogen, despite nitrogen being more electronegative. So I think a different explanation could me more accurate, but I’m not sure how to properly address it. I better keep reading.

In cyanide ion, there are two lone pairs – one on carbon, one on nitrogen. The lone pair on carbon is more nucleophilic because it is less tightly held (the atom is less electronegative than nitrogen). On all the examples I show that are negatively charged (eg BH4(-) ) there isn’t a lone pair to complicate questions of nucleophilicity.

This really helped for neutral covalent molecules. However, I’m having trouble applying this technique for molecules with an overall charge other than 0. For instance, in (ClO2)- , the formal charge of Cl should be 1. However, with your equation the charge should be 0. With the conventional equation, the charge is indeed 1.

I’d appreciate it if you replied sooner rather than later, as I do have a chemistry midterm on Friday. I’m quite confused with formal charges :)

Thanks for the study guide.

This method is wrong For CH3 , the valence eloctron is 4 , no : of bonds is 3 and no of non bonded electrons is 1 Then by this equation

F.C= 4-(1+3) = 0 but here it is given as +1

That analysis would be accurate for the methyl radical. However it fails for the methyl carbocation.

That example referred to the carbocation. For the methyl radical, the formal charge is indeed zero.

This was so helpful n the best explanation about the topic…

Thanks for the easy approach.

But when I used this formula it works. Thus #valence electrons_#lone pair__#1/2.bond pairs

Thanks for the easy approach. I have a problem in finding the FC on each O atom in ozone. Can you help me with that ASAP?

The FC on central atom would be +1 because [6-(2+3)] FC on O atom with coordinate bond would be: -1 because [6-(6+1)]. FC on O atom with double bond is: 0 because [6-(4+2)].

Hope I solved your question!

Thank u very much my exam is today and i wouldn’t pass without this information

AM REALLY LOST NOW ON THAT EXAMPLE OF CH3 CARBON # OF VALENCE ELECTRON=4 # OF BONDING=3 # OF UNSHARED=1

SO WHEN I CALCULATE

FORMAL CHARGE=(#OF VALENCE ELEC)+[(1/2#OF BOND)+(#OF UNSHARED)] FORMAL CHARGE=4+[(1/2*3)+1] =1.5

PLZ HELP IF AM MAKING MISTAKE

Should be 1/2 [# of bonding ELECTRONS] + # unshared. This gives you 4 – [3 – 1] = 0 for ch3 radical.

Should be for CH3(+), not the methyl radical •CH3 .

I am beryllium and i got offended!!!!!!……..LOL Just kidding…….BTW, I found this article very useful.Thanks!!!!!!!!!!

what does it means if we determine a molecule with zero charge ?

It’s neutral!

you said that non bonded electrons in carbon is 2, but how ? because i see it as only 1 because out of the 4 valence electrons in carbon, three are paired with hydrogen so it’s only 1 left

If the charge is -1, there must be an “extra” electron on carbon – this is why there’s a lone pair. If there was only one electron, it would be neutral.

This works! I would take your class with organic chemistry if you are a professor. I am taking chemistry 2 now. Organic is next. Thank you so much!

Thank you very very more for the simple explanation! Unbelievably easy and saves so much time!!!!!!

Thank you!!! this was awesome, I’m a junior in chemistry and this finally answered all my questions about formal charge :)

Glad it was helpful Haley!

If formal charges bear no resemblance to reality, what are their significance?

I hope the post doesn’t get interpreted as “formal charges have no significance”. If it does I will have to change some of the wording.

What I mean to get across is that formal charges assigned to atoms do not *always* accurately depict electron density on that atom, and one has to be careful.

In other words, formal charge and electron density are two different things and they do not always overlap.

Formal charge is a book-keeping device, where we count electrons and assign a full charge to one or more of the atoms on a molecule or ion. Electron density, on the other hand, is a measurement of where the electrons actually are (or aren’t) on a species, and those charges can be fractional or partial charges.

First of all, the charge itself is very real. The ions NH4+ , HO-, H3O+ and so on actually do bear a single charge. The thing to remember is that from a charge density perspective, that charge might be distributed over multiple atoms. Take an ion like H3O+, for example. H3O *does* bear a charge of +1,

However, if one thinks about where the electrons are in H3O+, one realizes that oxygen is more electronegative than hydrogen, and is actually “taking’ electrons from each hydrogen. If you look at an electron density map of H3O+ , one will see that the positive charge is distributed on the three hydrogens, and the oxygen actually bears a slight negative charge. There’s a nice map here.

http://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Aqueous_Solutions/The_hydronium_Ion

When we calculate formal charge for H3O+, we assign a charge of +1 to oxygen. This is for book keeping reasons. As a book-keeping device, it would be a royal pain to deal with fractions of charges like this. So that’s why we calculate formal charge and use it.

Sometimes it does accurately depict electron density. For example, in the hydroxide ion, HO- , the negative charge is almost all on the oxygen.

If you have a firm grasp of electronegativity then it becomes less confusing.

Does that help?

There are meny compounds which bears various structure among these which one is more stable or less energetic is it possible to predicu from the formal charge calculation?

Hey great explanation. I have a question though. Why is the FC commonly +/- 1? Could you give me an example when the FC is not +/- 1? Thanks.

Sure, try oxygen with no bonds and a full octet of electrons.

Great!i can use this for my exam!thanks!

Shouldn’t the formal charge of CH3 be -1? I was just wondering because in your example its +1 and in the chart its -1.

In the question.. its mentioned that CH3 without any lone pairs.. which means the valence would be 4 but there will not be any (2electrons) lone pairs left.. Hence it will be (4-)-(0+3)= 1

In CH3 i think FC on C should be -1 as carbon valency is 4 it has already bonded with 3 hydrogen atom one electron is left free on carbon to get bond with or share with one electron H hence, number of non bonded electrons lone pair of electrons is considered as 2. 4-(2+3) = -1. In your case if we take 0 than valency of c is not satisfied.

thank you for collaboration of formal charge

The answer to the question in the post above is “carbenes” – they have two substitutents, one pair of electrons, and an empty p orbital – so a total of four electrons “to itself”, making it neutral.

thank you for excellent explanation

Glad you found it useful Peter!

Very good explanation.I finally understood how to calculate the formal charge,was having some trouble with it.Thanks:)

Glad you found it helpful.

nice, concise explanation

sir the sheet posted by u is really very excellent.i m teacher of chemistry in india for pre engineering test.if u send me complete flow chart of chemistry i will great full for u

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Module 7: Chemical Bonding and Molecular Geometry

Formal charges and resonance, learning outcomes.

  • Compute formal charges for atoms in any Lewis structure
  • Use formal charges to identify the most reasonable Lewis structure for a given molecule
  • Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule

In the previous section, we discussed how to write Lewis structures for molecules and polyatomic ions. As we have seen, however, in some cases, there is seemingly more than one valid structure for a molecule. We can use the concept of formal charges to help us predict the most appropriate Lewis structure when more than one is reasonable.

Calculating Formal Charge

The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.

Thus, we calculate formal charge as follows:

[latex]\text{formal charge}=\text{# valence shell electrons (free atom)}-\text{# lone pair electrons}-\dfrac{1}{2}\text{# bonding electrons}[/latex]

We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.

We must remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.

Example 1: Calculating Formal Charge from Lewis Structures

Assign formal charges to each atom in the interhalogen ion [latex]{\text{ICl}}_{4}^{-}.[/latex]

A Lewis structure is shown. An iodine atom with two lone pairs of electrons is single bonded to four chlorine atoms, each of which has three lone pairs of electrons. Brackets surround the structure and there is a superscripted negative sign.

  • We assign lone pairs of electrons to their atoms . Each Cl atom now has seven electrons assigned to it, and the I atom has eight.
  • Subtract this number from the number of valence electrons for the neutral atom: I: 7 – 8 = –1 Cl: 7 – 7 = 0 The sum of the formal charges of all the atoms equals –1, which is identical to the charge of the ion (–1).

Check Your Learning

Calculate the formal charge for each atom in the carbon monoxide molecule:

A Lewis structure is shown. A carbon atom with one lone pair of electrons is triple bonded to an oxygen with one lone pair of electrons.

Example 2: Calculating Formal Charge from Lewis Structures

Assign formal charges to each atom in the interhalogen molecule BrCl 3 .

A Lewis structure is shown. A bromine atom with two lone pairs of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.

  • Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.
  • Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge:Br: 7 – 7 = 0Cl: 7 – 7 = 0All atoms in BrCl 3 have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.

Determine the formal charge for each atom in NCl 3 .

A Lewis structure is shown. A nitrogen atom with one lone pair of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.

Using Formal Charge to Predict Molecular Structure

The arrangement of atoms in a molecule or ion is called its molecular structure . In many cases, following the steps for writing Lewis structures may lead to more than one possible molecular structure—different multiple bond and lone-pair electron placements or different arrangements of atoms, for instance. A few guidelines involving formal charge can be helpful in deciding which of the possible structures is most likely for a particular molecule or ion:

  • A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.
  • If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.
  • Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.
  • When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

To see how these guidelines apply, let us consider some possible structures for carbon dioxide, CO 2 . We know from our previous discussion that the less electronegative atom typically occupies the central position, but formal charges allow us to understand why this occurs. We can draw three possibilities for the structure: carbon in the center and double bonds, carbon in the center with a single and triple bond, and oxygen in the center with double bonds:

Three Lewis structures are shown. The left and right structures show a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons. The center structure shows a carbon atom that is triple bonded to an oxygen atom with one lone pair of electrons and single bonded to an oxygen atom with three lone pairs of electrons. The third structure shows an oxygen atom double bonded to another oxygen atom with to lone pairs of electrons. The first oxygen atom is also double bonded to a carbon atom with two lone pairs of electrons.

Comparing the three formal charges, we can definitively identify the structure on the left as preferable because it has only formal charges of zero (Guideline 1).

As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular structures: CNS – , NCS – , or CSN – . The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:

Two rows of structures and numbers are shown. The top row is labeled, “Structure” and depicts three Lewis structures and the bottom row is labeled, “Formal charge.” The left structure shows a carbon atom double bonded to a nitrogen atom with two lone electron pairs on one side and double bonded to a sulfur atom with two lone electron pairs on the other. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative one, zero, and zero. The middle structure shows a carbon atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to a sulfur atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive one, and zero. The right structure shows a carbon atom with two lone electron pairs double bonded to a sulfur atom that is double bonded to a nitrogen atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive two, and one.

Note that the sum of the formal charges in each case is equal to the charge of the ion (–1). However, the first arrangement of atoms is preferred because it has the lowest number of atoms with nonzero formal charges (Guideline 2). Also, it places the least electronegative atom in the center, and the negative charge on the more electronegative element (Guideline 4).

Example 3: Using Formal Charge to Determine Molecular Structure

Nitrous oxide, N 2 O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth. Which is the likely structure for nitrous oxide?

Two Lewis structures are shown with the word “or” in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen that is double bonded to an oxygen with two lone pairs of electrons. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons.

Determining formal charge yields the following:

Two Lewis structures are shown with the word “or” in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons. The numbers negative one, positive one, and zero are written above this structure. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons. The numbers negative one, positive two, and negative one are written above this structure.

The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:

A Lewis structure is shown. A nitrogen atom with two lone pairs of electrons is double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons.

The number of atoms with formal charges are minimized (Guideline 2), and there is no formal charge larger than one (Guideline 2). This is again consistent with the preference for having the less electronegative atom in the central position.

Which is the most likely molecular structure for the nitrite [latex]\left({\text{NO}}_{2}^{-}\right)[/latex] ion?

Two Lewis structures are shown with the word “or” written between them. The left structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign.

If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. Experiments show, however, that both N–O bonds in [latex]{\text{NO}}_{2}^{-}[/latex] have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for [latex]{\text{NO}}_{2}^{-}[/latex] in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of resonance : if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in [latex]{\text{NO}}_{2}^{-}[/latex] is the average of a double bond and a single bond. We call the individual Lewis structures resonance forms . The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms. Thus, the electronic structure of the [latex]{\text{NO}}_{2}^{-}[/latex] ion is shown as:

Two Lewis structures are shown with a double sided arrow between them. The left structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.

We should remember that a molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn, because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).

Three Lewis structures are shown with double headed arrows in between. Each structure is surrounded by brackets, and each has a superscripted two negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to two of these oxygen atoms, each of which has three lone pairs of electrons, and double bonded to the third, which has two lone pairs of electrons. The double bond is located between the lower left oxygen atom and the carbon atom. The central and right structures are the same as the first, but the position of the double bonded oxygen has moved to the lower right oxygen in the central structure and to the top oxygen in the right structure.

One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C–O bonds are exactly the same.

You can view the transcript for “Resonance” here (opens in new window) .

Key Concepts and Summary

In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).

Key Equations

  • [latex]\text{formal charge}=\text{# valence shell electrons (free atom)}-\text{# one pair electrons}-\dfrac{1}{2}\text{# bonding electrons}[/latex]
  • selenium dioxide, OSeO
  • nitrate ion, [latex]{\text{NO}}_{3}^{-}[/latex]
  • nitric acid, HNO 3 (N is bonded to an OH group and two O atoms)

A Lewis structure shows a hexagonal ring composed of six carbon atoms. They form single bonds to each another and single bonds to one hydrogen atom each.

  • sulfur dioxide, SO 2 
  • carbonate ion, [latex]{\text{CO}}_{3}^{2-}[/latex]
  • hydrogen carbonate ion, [latex]{\text{HCO}}_{3}^{-}[/latex] (C is bonded to an OH group and two O atoms)

A Lewis structure depicts a hexagonal ring composed of five carbon atoms and one nitrogen atom. Each carbon atom is single bonded to a hydrogen atom.

  • Write the resonance forms of ozone, O 3 , the component of the upper atmosphere that protects the Earth from ultraviolet radiation.
  • Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, [latex]{\text{NO}}_{\text{2}}^{-}.[/latex]

Two Lewis structures are shown with a double headed arrow in between. The left structure shows a carbon atom single bonded to three hydrogen atoms and a second carbon atom. The second carbon is single bonded to two oxygen atoms. One of the oxygen atoms is single bonded to a hydrogen atom. The right structure, surrounded by brackets and with a superscripted negative sign, depicts a carbon atom single bonded to three hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to two oxygen atoms.

  • Toothpastes containing sodium hydrogen carbonate (sodium bicarbonate) and hydrogen peroxide are widely used. Write Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule, with resonance forms where appropriate.
  • [latex]{\text{SO}}_{4}^{2-}[/latex]
  • [latex]{\text{O}}_{2}^{2-}[/latex] (e) H 2 O 2
  • Calculate the formal charge of chlorine in the molecules Cl 2 , BeCl 2 , and ClF 5 .
  • [latex]{\text{BF}}_{4}^{-}[/latex]
  • [latex]{\text{SnCl}}_{3}^{-}[/latex]
  • [latex]{\text{PO}}_{4}^{\text{3-}}[/latex]
  • [latex]{\text{NO}}_{2}^{-}[/latex]
  • [latex]{\text{NO}}_{3}^{-}[/latex]
  • Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in nitrosyl chloride: ClNO or ClON?
  • Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?
  • Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in sulfur dioxide: OSO or SOO?
  • Draw the structure of hydroxylamine, H 3 NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?
  • Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.

Two Lewis structures are shown, with the word “or” in between. The left structure shows a nitrogen atom single bonded to an oxygen atom with three lone pairs of electrons. It is also single bonded to a hydrogen atom and double bonded to an oxygen atom with two lone pairs of electrons. The right structure shows a hydrogen atom single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to a nitrogen atom which is double bonded to an oxygen atom with two lone pairs of electrons.

  • Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H 2 SO 4 , which has two oxygen atoms and two OH groups bonded to the sulfur.

2. The resonance forms are as follows:

Two Lewis structures are shown with a double-headed arrow in between. The left structure shows a sulfur atom with a lone pair of electrons single bonded to the left to an oxygen atom with three lone pairs of electrons. The sulfur atom is also double bonded on the right to an oxygen atom with two lone pairs of electrons. The right structure depicts the same atoms, but this time the double bond is between the left oxygen and the sulfur atom. The lone pairs of electrons have also shifted to account for the change of bond types. The sulfur atom in the right structures, also has a third electron dot below it.

6. The Lewis structures are as follows:

This structure shows a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons.

CO has the strongest carbon-oxygen bond, because there are is a triple bond joining C and O. CO 2 has double bonds, and carbonate has 1.3 bonds.

12. Draw all possible resonance structures for each of the compounds below. Determine the formal charge on each atom in each of the resonance structures:

Two Lewis structures are shown with a double-headed arrow in between. The left structure shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, “( 0 ), ( positive 1 ), ( negative 1 ).” The phrase, “Formal charge,” and a right-facing arrow lie to the left of this structure. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, “( negative 1 ), ( positive 1 ), ( 0 ).”

The structure with formal charges of 0 is the most stable and would therefore be the correct arrangement of atoms.

A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to two hydrogen atoms and an oxygen atom which has two lone pairs of electrons. The oxygen atom is single bonded to a hydrogen atom.

18. There are 19.7 g N and 80.3 g F in a 100.0-g sample:

[latex]\begin{array}{l}\dfrac{19.7\text{g}}{14.0067\text{ g}{\text{ mol}}^{-1}}=1.406\text{ mol}\\ \dfrac{1.406\text{ mol}}{1.406\text{ mol}}=1\text{ N}\\ \dfrac{80.3\text{ g}}{18.9984\text{ g}{\text{ mol}}^{-1}}=4.2267\text{ mol}\\ \dfrac{4.2267\text{ mol}}{1.406\text{ mol}}=3\text{ F}\end{array}[/latex]

The empirical formula is NF 3 and its molar mass is 71.00 g/mol, which is consistent with the stated molar mass.

  • Oxidation states: N = +3, F = –1.

A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to three fluorine atoms, each with three lone pairs of electrons.

formal charge:  charge that would result on an atom by taking the number of valence electrons on the neutral atom and subtracting the nonbonding electrons and the number of bonds (one-half of the bonding electrons)

molecular structure:  arrangement of atoms in a molecule or ion

resonance:  situation in which one Lewis structure is insufficient to describe the bonding in a molecule and the average of multiple structures is observed

resonance forms:  two or more Lewis structures that have the same arrangement of atoms but different arrangements of electrons

resonance hybrid:  average of the resonance forms shown by the individual Lewis structures

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Course: organic chemistry   >   unit 2.

  • Comparing formal charges to oxidation states
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Formal charge on nitrogen

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Learn Chemistry Online | ChemistryScore

How to calculate formal charge

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How to calculate formal charge Examples

ot all atoms within a neutral molecule need be neutral. An atom can have the following charges: positive , negative , or neutral , depending on the electron distribution. This is often useful for understanding or predicting reactivity. Identifying formal charges helps you keep track of the electrons.

The formal charge is the charge on the atom in the molecule. The term “formal” means that this charge is not necessarily on the presented atom because in some cases, it is also prevalent on other atoms present in the molecule. It is actually spread out through the other atoms and is not only on the one atom. Identifying a formal charge involves:

  • Determining the appropriate number of valence electrons for an atom – This can be accomplished by inspecting the periodic table. The group number indicates the appropriate number of valence electrons for each atom
  • Determining whether the atom exhibits the appropriate number of electrons – In the Lewis structure, determine whether some of the atoms show an unexpected number of electrons

The formal charge on an atom can be calculated using the following mathematical equation.

formal-charge

Lewis structures also show how atoms in the molecule are bonded. They can be drawn as lines (bonds) or dots (electrons). One line corresponds to two electrons . The nonbonding electrons, on the other hand, are the unshared electrons and these are shown as dots. One dot is equal to one nonbonding electron. The valence electrons are the electrons in the outermost shell of the atom.

formal-charge-2

CH 4 , methane

CH4-Methane

A number of non-bonding electrons: 0 for both H and C

[ Formal charge ] H = 1 – (1/2) × 2 – 0 = 0 ⇒ This applies to each hydrogen. These hydrogens are all zero.

[ Formal charge ] C = 4 – (1/2) × 8 – 0 = 0

⇒ This molecule is neutral .

CH 3 + , methyl cation

CH3-Methyl-cation

[ Formal charge ] H = 1 – (1/2) × 2 – 0 = 0 ⇒ This applies to each hydrogen. These hydrogens are all zero. [ Formal charge ] C = 4 – (1/2) × 6 – 0 = 4 – 3 – 0 = +1

⇒ This is a cation .

CH 3 – , methyl cation

CH-3-methyl-cation

A number of non-bonding electrons: 0 for H, 2 for C

[ Formal charge ] C = 4 – (1/2) × 6 – 2 = 4 – 3 – 2 = -1

⇒ This is a anion .

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Calculations and examples, how to calculate formal change.

Refresher: Each atom on the periodic table has a set number of valence electrons. The number of valence electrons increases across the periodic table. When atoms form covalent bonds, one pair of electrons is shared between the atoms for every single bond. Two pairs of electrons are shared in a double bond and three pairs are shared in a triple bond.

Definition of formal charge

Formal charge is the charge placed on each atom in a molecule, ignoring relative electronegativities. It is a result of the electrons associated with each atom in the molecule. Even if a molecule is overall neutral, a formal charge can be associated with the atoms in the molecule (which would add up to result in an overall zero charge). Knowing how electrons are distributed in a molecule and the formal charge helps us to predict how the molecule will react.

calculate formal charge of no2

Main Takeaways

  • You can calculate formal charge by determining the amount of valence electrons an atom should have and subtracting the amount of electrons it has in the molecule of interest.
  • Each lone pair of electrons is two electrons that the atom has and each bond contributes one electron.
  • All of the formal charges should add up to the charge on the overall molecule.

Periodic Tables

What is the formal charge of N O − 2 ?

Formal charge on atom in a molecule= (total no. of valence electrons in the free atom)- (total no. of non-bonding electrons)-1/2 (total no. of bonding electrons) formal charge on double bonded o atom = 6 – 4 – 0.5 ( 4 ) = 0 formal charge on single bonded o atom = 6 – 6 – 0.5 ( 2 ) = − 1.

What is the formal charge of C in CO?

What is lowest energy structure, in terms of formal charge?

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7.4: Formal Charges and Resonance

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  Learning Objectives

  • Compute formal charges for atoms in any Lewis structure
  • Use formal charges to identify the most reasonable Lewis structure for a given molecule
  • Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule

Previously, we discussed how to write Lewis structures for molecules and polyatomic ions. In some cases, however, there is seemingly more than one valid structure for a molecule. We can use the concept of formal charges to help us predict the most appropriate Lewis structure when more than one is reasonable.

Calculating Formal Charge

The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.

Thus, we calculate formal charge as follows:

\[\textrm{formal charge = # valence shell electrons (free atom) − # lone pair electrons − }\dfrac{1}{2}\textrm{ # bonding electrons} \nonumber \]

We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.

We must remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.

Example \(\PageIndex{1}\): Calculating Formal Charge from Lewis Structures

Assign formal charges to each atom in the interhalogen ion \(\ce{ICl4-}\).

We divide the bonding electron pairs equally for all \(\ce{I–Cl}\) bonds:

imageedit_2_8215662475.png

We assign lone pairs of electrons to their atoms. Each Cl atom now has seven electrons assigned to it, and the I atom has eight.

Subtract this number from the number of valence electrons for the neutral atom:

  • I: 7 – 8 = –1
  • Cl: 7 – 7 = 0

The sum of the formal charges of all the atoms equals –1, which is identical to the charge of the ion (–1).

Exercise \(\PageIndex{1}\)

Calculate the formal charge for each atom in the carbon monoxide molecule:

imageedit_9_8783034466.png

C −1, O +1

Example \(\PageIndex{2}\): Calculating Formal Charge from Lewis Structures

Assign formal charges to each atom in the interhalogen molecule \(\ce{BrCl3}\).

Assign one of the electrons in each Br–Cl bond to the Br atom and one to the Cl atom in that bond:

imageedit_6_8383188586.png

Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.

Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge:

  • Br: 7 – 7 = 0

All atoms in \(\ce{BrCl3}\) have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.

Exercise \(\PageIndex{2}\)

Determine the formal charge for each atom in \(\ce{NCl3}\).

N: 0; all three Cl atoms: 0

imageedit_28_5766396641.png

Using Formal Charge to Predict Molecular Structure

The arrangement of atoms in a molecule or ion is called its molecular structure . In many cases, following the steps for writing Lewis structures may lead to more than one possible molecular structure—different multiple bond and lone-pair electron placements or different arrangements of atoms, for instance. A few guidelines involving formal charge can be helpful in deciding which of the possible structures is most likely for a particular molecule or ion.

Predicting Molecular Structure Guidelines

  • A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.
  • If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.
  • Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.
  • When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

To see how these guidelines apply, let us consider some possible structures for carbon dioxide, \(\ce{CO2}\). We know from our previous discussion that the less electronegative atom typically occupies the central position, but formal charges allow us to understand why this occurs. We can draw three possibilities for the structure: carbon in the center and double bonds, carbon in the center with a single and triple bond, and oxygen in the center with double bonds:

Comparing the three formal charges, we can definitively identify the structure on the left as preferable because it has only formal charges of zero (Guideline 1).

As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular structures: \(\ce{CNS^{–}}\), \(\ce{NCS^{–}}\), or \(\ce{CSN^{–}}\). The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:

Note that the sum of the formal charges in each case is equal to the charge of the ion (–1). However, the first arrangement of atoms is preferred because it has the lowest number of atoms with nonzero formal charges (Guideline 2). Also, it places the least electronegative atom in the center, and the negative charge on the more electronegative element (Guideline 4).

Example \(\PageIndex{3}\): Using Formal Charge to Determine Molecular Structure

Nitrous oxide, N 2 O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth. Which is the likely structure for nitrous oxide?

imageedit_13_7711983386.png

Solution Determining formal charge yields the following:

imageedit_16_6226662442.png

The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:

imageedit_20_2309205655.png

The number of atoms with formal charges are minimized (Guideline 2), and there is no formal charge larger than one (Guideline 2). This is again consistent with the preference for having the less electronegative atom in the central position.

Exercise \(\PageIndex{3}\)

Which is the most likely molecular structure for the nitrite (\(\ce{NO2-}\)) ion?

imageedit_24_5491927072.png

\(\ce{ONO^{–}}\)

You may have noticed that the nitrite anion in Example \(\PageIndex{3}\) can have two possible structures with the atoms in the same positions. The electrons involved in the N–O double bond, however, are in different positions:

If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. Experiments show, however, that both N–O bonds in \(\ce{NO2-}\) have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for \(\ce{NO2-}\) in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of resonance : if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in \(\ce{NO2-}\) is the average of a double bond and a single bond. We call the individual Lewis structures resonance forms . The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms. Thus, the electronic structure of the \(\ce{NO2-}\) ion is shown as:

We should remember that a molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).

The carbonate anion, \(\ce{CO3^2-}\), provides a second example of resonance:

One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C–O bonds are exactly the same.

In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).

Key Equations

  • \(\textrm{formal charge = # valence shell electrons (free atom) − # one pair electrons − }\dfrac{1}{2}\textrm{ # bonding electrons}\)

IMAGES

  1. Calculating NO2- Formal Charges: Calculating Formal Charges for the Nitrite Ion

    calculate formal charge of no2

  2. NO2- Formal charge, How to calculate it with images?

    calculate formal charge of no2

  3. NO2- Formal charge, How to calculate it with images?

    calculate formal charge of no2

  4. NO2- Formal charge, How to calculate it with images?

    calculate formal charge of no2

  5. How to Calculate the Formal Charges for NO2 (Nitrogen dioxide)

    calculate formal charge of no2

  6. NO2- Formal charge, How to calculate it with images?

    calculate formal charge of no2

VIDEO

  1. How to calculate Formal charge |class-XI

  2. How you can calculate Formal Charge on atoms. 📚

  3. Formal charge of CO

  4. Formal charge calculation |formal charge on Co2

  5. General Chemistry

  6. Trick to calculate Formal charge || Easiest way to find formal charge of an atom

COMMENTS

  1. How to Calculate the Formal Charges for NO2 (Nitrogen dioxide)

    In order to calculate the formal charges for NO2 we'll use the equation:Formal charge = [# of valence electrons] - [nonbonding val electrons] - [bonding elec...

  2. 8.4: Formal Charge

    The sum of the formal charges of all the atoms in a neutral molecule equals zero; The sum of the formal charges of all the atoms in an ion equals the charge of the ion. Uses of Formal Charges. Formal charges can help identify the more important resonance structures, that is, hitherto we have treated all resonance structures as equal, but this ...

  3. How to count formal charge in NO2

    Formal charge of NO 2-:-Structure of NO 2-is: Step 1: Formal charge of Nitrogen . Here Nitrogen is the free atom and the number of valence electrons of it is 5. Number of non-bonding electrons is 2 and bonding electrons are 6. ∴ Formal charge of Nitrogen is . FC = V − N − B 2 ⇒ FC = 5-2-(6 2) ⇒ FC = 5-5 ⇒ FC = 0. Step 2: Formal ...

  4. How to count formal charge in NO2?

    And so ""^(-)O-stackrel(ddot)N=O...around each atom from your left to right there are 9, 7, and 8 electrons respectively leading to formal charges of -1, 0, and 0. Of course we can distribute the negative charge over the two oxygen centres by resonance, and so /_O-N-O < 120^@, i.e. the nitrogen lone pair, which is closer to the nitrogen atom ...

  5. NO2 Lewis Structure: Drawings, Hybridization, Shape, Charges, Pairs

    Calculation of Formal Charges in NO2. To calculate the formal charges in a NO2 molecule, we need to follow a step-by-step process:. Determine the total number of valence electrons in the molecule. For NO2, nitrogen (N) contributes 5 valence electrons, and each oxygen (O) contributes 6 valence electrons, giving us a total of 5 + 2(6) = 17 valence electrons.

  6. 4.3: Formal Charge and Oxidation State

    The C atom has gained four electrons, giving it a negative charge and hence an oxidation number of - 4: C−4H+1 4 (4.3.3) (4.3.3) C − 4 H +1 4. c) In NaCl each Na atom has lost an electron to form an Na + ion, and each Cl atom has gained an electron to form Cl -.

  7. Formal Charges: Calculating Formal Charge

    A step-by-step description on how to calculate formal charges. Formal charges are important because they allow us to predict which Lewis structure is the mo...

  8. 7.4: Formal Charges and Resonance

    Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge: Br: 7 - 7 = 0. Cl: 7 - 7 = 0. All atoms in BrCl3 BrCl 3 have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.

  9. Formal charge (video)

    Formal charge. The formal charge of an atom in a molecule is the charge that would reside on the atom if all of the bonding electrons were shared equally. We can calculate an atom's formal charge using the equation FC = VE - [LPE - ½ (BE)], where VE = the number of valence electrons on the free atom, LPE = the number of lone pair electrons on ...

  10. 7.4 Formal Charges and Resonance

    Calculate the formal charge for each atom in the carbon monoxide molecule: Answer: C −1, O +1. Example 7.7. Calculating Formal Charge from Lewis Structures Assign formal charges to each atom in the interhalogen molecule BrCl 3. Solution. Step 1.

  11. How To Calculate Formal Charge

    How To Calculate Formal Charge. To calculate the formal charge of an atom, we start by:. evaluating the number of valence electrons (VE) the neutral atom has (e.g. 3 for boron, 4 for carbon, 5 for nitrogen, and so on). (note: this is also equivalent to the effective nuclear charge Z eff, the number of protons that an electron in the valence orbital "sees" due to screening by inner-shell ...

  12. Formal Charges and Resonance

    Thus, we calculate formal charge as follows: formal charge= # valence shell electrons (free atom)−# lone pair electrons− 1 2# bonding electrons formal charge = # valence shell electrons (free atom) − # lone pair electrons − 1 2 # bonding electrons. We can double-check formal charge calculations by determining the sum of the formal ...

  13. Formal charge on nitrogen (video)

    The formal charge on N is usually -1 for an anion, 0 for a neutral compound, and +1 in cations. A nitrogen atom with a formal charge of -3 would correspond to a nitride ion, N³⁻, which is strongly basic in aqueous solution. A nitrogen atom surrounded by four hydrogen atoms is the ammonium radical (NH4+).

  14. How to calculate formal charge

    The formal charge on an atom can be calculated using the following mathematical equation.. Lewis structures also show how atoms in the molecule are bonded. They can be drawn as lines (bonds) or dots (electrons).One line corresponds to two electrons.The nonbonding electrons, on the other hand, are the unshared electrons and these are shown as dots.

  15. How to calculate formal change

    You can calculate formal charge by determining the amount of valence electrons an atom should have and subtracting the amount of electrons it has in the molecule of interest. Each lone pair of electrons is two electrons that the atom has and each bond contributes one electron. All of the formal charges should add up to the charge on the overall ...

  16. Calculate the formal charges for each of the oxygen atoms within

    Calculate the formal charges for each of the oxygen atoms within the nitrite ion, NO2-.

  17. 2.2: Formal Charges

    Using Equation 2.2.1 2.2.1 to calculate the formal charge on hydrogen, we obtain. FC(H) = (1 valence electrons) − (0 lone pair electrons) − 1 2(2 bonding electrons) = 0 F C ( H) = ( 1 valence electrons) − ( 0 lone pair electrons) − 1 2 ( 2 bonding electrons) = 0. The sum of the formal charges of each atom must be equal to the overall ...

  18. Solved CALCULATE THE FORMAL CHARGES OF ALL ATOMS IN NH3,

    Question: CALCULATE THE FORMAL CHARGES OF ALL ATOMS IN NH3, NO2-, & NO3-. Here's the best way to solve it. equation FOR CALCULATING FORMAL CHARGES:formal cha ….

  19. What is the formal charge of { NO }_{ 2 }^{

    Formal charge on atom in a molecule= (total no. of valence electrons in the free atom)- (total no. of non-bonding electrons)-1/2 (total no. of bonding electrons) Formal charge on double bonded O atom = 6 - 4 - 0.5 ( 4 ) = 0

  20. Solved 2. Calculate the formal charges of all atoms in

    2. Calculate the formal charges of all atoms in NH3,NO2−, and NO3−. 3. Explain the origin of the octet rule. Which principle is it based on? 4. Describe in your own words polar and nonpolar bonds. 5. A chemical compound has polar bonds, but it is a nonpolar molecule.

  21. 2.3: Formal Charges

    The formal charge of each atom in a molecule can be calculated using the following equation: Formal Charge = (# of valence electrons in free atom) - (# of lone-pair electrons) - (1/2 # of bond pair electrons) Eqn. 2.3.1. To illustrate this method, let's calculate the formal charge on the atoms in ammonia (NH 3) whose Lewis structure is as ...

  22. Formal Charge

    The formal charge of any atom in a molecule can be calculated by the following equation: FC = V − N − B 2 (1) (1) F C = V − N − B 2. where V is the number of valence electrons of the neutral atom in isolation (in its ground state); N is the number of non-bonding valence electrons on this atom in the molecule; and B is the total number ...

  23. 7.4: Formal Charges and Resonance

    Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge: Br: 7 - 7 = 0. Cl: 7 - 7 = 0. All atoms in BrCl3 BrCl 3 have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.