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.
Ch.8 - Basic Concepts of Chemical Bonding
Chapter 8, Problem 91b
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?
Verified step by step guidance1
Step 1: For the hypochlorite ion (ClO^-), start by counting the total number of valence electrons. Chlorine has 7 valence electrons, oxygen has 6, and the negative charge adds 1 more electron, totaling 14 electrons.
Step 2: Draw the skeletal structure for ClO^-. Place Cl and O next to each other and connect them with a single bond. Distribute the remaining electrons to satisfy the octet rule, starting with the more electronegative atom, oxygen.
Step 3: For the perchlorate ion (ClO4^-), count the total number of valence electrons. Chlorine has 7 valence electrons, each oxygen has 6, and the negative charge adds 1 more electron, totaling 32 electrons.
Step 4: Draw the skeletal structure for ClO4^-. Place Cl in the center and connect it to four O atoms with single bonds. Distribute the remaining electrons to satisfy the octet rule for each oxygen atom.
Step 5: To find the formal charge of Cl in ClO4^-, use the formula: Formal Charge = (Valence electrons of Cl) - (Non-bonding electrons on Cl) - (Bonding electrons/2). Assume all Cl—O bonds are single bonds.
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Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
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 help visualize the arrangement of electrons and the connectivity of atoms, which is crucial for understanding molecular geometry and reactivity. In the case of ions like ClO- and ClO4-, drawing the Lewis structure allows us to see how the chlorine and oxygen atoms are bonded and the distribution of electrons.
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Formal Charge
Formal charge is a concept used to determine the charge of an atom in a molecule or ion, calculated by comparing the number of valence electrons in the free atom to the number of electrons assigned to it in the Lewis structure. It helps assess the stability of a molecule and the distribution of charge. For the perchlorate ion, calculating the formal charge of chlorine involves considering its valence electrons and the bonds it forms with oxygen.
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Ionic and Covalent Bonds
Ionic and covalent bonds are two primary types of chemical bonds. Ionic bonds occur when electrons are transferred from one atom to another, resulting in charged ions, while covalent bonds involve the sharing of electrons between atoms. In the context of the hypochlorite and perchlorate ions, understanding these bonding types is essential for analyzing their structures and properties, particularly when considering the nature of the Cl—O bonds.
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Related Practice
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
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).
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?
Textbook Question
The following three Lewis structures can be drawn for N2O:
(b) The N—N bond length in N2O is 1.12 Å, slightly longer than a typical N≡N bond; and the N—O bond length is 1.19 Å, slightly shorter than a typical N═O bond (see Table 8.4). Based on these data, which resonance structure best represents N2O?