Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in O 3 (b) phosphorus in PF 6 - (c) nitrogen in NO 2 (d) iodine in ICl 3 (e) chlorine in HClO 4 (hydrogen is bonded to O).

Hey everyone. So here says, shown below is the lewis structure of bromate ion. First it says, what is the formal charge of the bro mean, adam? Alright, so we're gonna say that formal charge sort of the browning formal charge equals the the group number of the element which is seven, since browning is in group seven A minus the number of bond it's making, which is five plus each individual non bonding electron. So the formal charge of bromine here would be zero. Alright, so now let's do the formal charge of each oxygen A, B, and C. Since A and B are the same, we'll group them together. So the oxygen that are double bonded with two lone pairs, we're gonna say look at one of the oxygen's let's look at A. So it is in group six a. We see making two bonds and it has four electrons not bonding, It would be zero. So so far we see that this is zero, this is zero. If oxygen A zero then oxygen B is zero. Now let's look at oxygen. See, It's formal charge, it is in group six a. We see making one bond and it has six individual electrons not bonding. So it's formal charges -1. Okay, so that's what we'd say for each of them. Now bromate, overall is bureau three minus one. It has a minus one charge. If you add up all the formal charges together, it's supposed to give you back the charge of the overall ion, which it does. Alright, so those will be our final answers

Ch.8 - Basic Concepts of Chemical Bonding

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Brown 14th Edition

Ch.8 - Basic Concepts of Chemical Bonding

Problem 90

Chapter 8, Problem 90

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in O₃ (b) phosphorus in PF₆^- (c) nitrogen in NO₂ (d) iodine in ICl₃ (e) chlorine in HClO₄ (hydrogen is bonded to O).

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To calculate the formal charge on an atom, use the formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (Bonding Electrons / 2).

For (a) the central oxygen atom in O3: Determine the number of valence electrons for oxygen, count the non-bonding electrons (lone pairs) on the central oxygen, and count the bonding electrons (shared in bonds) around it.

For (b) phosphorus in PF6-: Identify the valence electrons for phosphorus, count the non-bonding electrons, and the bonding electrons shared with fluorine atoms. Remember to account for the negative charge on the ion.

For (c) nitrogen in NO2: Find the valence electrons for nitrogen, count the non-bonding electrons, and the bonding electrons shared with oxygen atoms. Consider the resonance structures if applicable.

For (d) iodine in ICl3 and (e) chlorine in HClO4: Follow the same process by identifying the valence electrons, counting non-bonding electrons, and bonding electrons. For HClO4, note that hydrogen is bonded to oxygen, affecting the electron count around chlorine.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Formal Charge

Formal charge is a theoretical charge assigned to an atom in a molecule, calculated based on the number of valence electrons, the number of non-bonding electrons, and half the number of bonding electrons. It helps in determining the most stable structure of a molecule by minimizing the formal charges across the atoms. The formula for calculating formal charge is: Formal Charge = Valence Electrons - (Non-bonding Electrons + 1/2 Bonding Electrons).

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist. They are essential for visualizing the arrangement of electrons and understanding how atoms share or transfer electrons to form bonds. Drawing accurate Lewis structures is crucial for calculating formal charges, as it provides a clear picture of the electron distribution in the molecule.

Resonance Structures

Resonance structures are different Lewis structures that represent the same molecule, where the arrangement of electrons differs but the arrangement of atoms remains the same. They are used to depict delocalized electrons that cannot be represented by a single Lewis structure. Understanding resonance is important for accurately determining formal charges, as the actual charge distribution in a molecule may be an average of multiple resonance forms.

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Resonance Structures

Related Practice

Textbook Question

A major challenge in implementing the 'hydrogen economy' is finding a safe, lightweight, and compact way of storing hydrogen for use as a fuel. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. For example, NaAlH₄ can release 5.6% of its mass as H2 upon decomposing to NaH(s), Al(s), and H₂(g). NaAlH₄ possesses both covalent bonds, which hold polyatomic anions together, and ionic bonds. (d) What is the formal charge on hydrogen in the polyatomic ion?

Textbook Question

Although I₃^- is a known ion, F₃^- is not. (b) One of your classmates says that F₃^- does not exist because F is too electronegative to make bonds with another atom. Give an example that proves your classmate is wrong.

Textbook Question

Although I₃^- is a known ion, F₃^- is not. (c) Another classmate says F₃^- does not exist because it would violate the octet rule. Is this classmate possibly correct?

Textbook Question

The hypochlorite ion, ClO^-, is the active ingredient in bleach. The perchlorate ion, ClO₄^-, is a main component of rocket propellants. Draw Lewis structures for both ions. (b) What is the formal charge of Cl in the perchlorate ion, assuming the Cl—O bonds are all single bonds?

Textbook Question

The hypochlorite ion, ClO^-, is the active ingredient in bleach. The perchlorate ion, ClO₄^-, is a main component of rocket propellants. Draw Lewis structures for both ions. (c) What is the oxidation number of Cl in the hypochlorite ion? (d) What is the oxidation number of Cl in the perchlorate ion, assuming the Cl—O bonds are all single bonds?

Textbook Question

The following three Lewis structures can be drawn for N2O:

(a) Using formal charges, which of these three resonance forms is likely to be the most important?

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in O3 (b) phosphorus in PF6- (c) nitrogen in NO2 (d) iodine in ICl3 (e) chlorine in HClO4 (hydrogen is bonded to O).

Verified video answer for a similar problem:

Key Concepts

Formal Charge

Lewis Structures

Resonance Structures

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in O₃ (b) phosphorus in PF₆^- (c) nitrogen in NO₂ (d) iodine in ICl₃ (e) chlorine in HClO₄ (hydrogen is bonded to O).