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Chapter 1: Problem 9

Draw Lewis structures for the following charged species. In each case, the charge is shown closest to the charged atom. (a) \(-\mathrm{OH}\) (b) \(-\mathrm{BH}_{4}\) (c) \({ }^{+} \mathrm{NH}_{4}\) (d) \({ }^{-} \mathrm{Cl}\) (e) \({ }^{+} \mathrm{CH}_{3}\) (f) \({ }^{+} \mathrm{OH}_{3}\) (g) \({ }^{+} \mathrm{NO}_{2}\)

Short Answer

Expert verified
Lewis structures: (a) :O:--H:(b) H--B--H | | H H (c) H--N--H | H(d) :Cl:(e) H--C--H | (f) H | O--H | H (g) O(+)==N--O(-)

Step by step solution

01

- Identify Valence Electrons

For each species, count the number of valence electrons from each atom including the additional charges.
02

- Draw Skeleton Structures

Draw the skeletal structures of each species by connecting atoms with single bonds. Place the central atom accordingly.
03

- Distribute Electrons

Distribute the remaining valence electrons around the atoms to satisfy the octet rule (or duet rule for hydrogen), starting with the outer atoms.
04

- Address Formal Charges

Adjust the electron distribution if necessary to minimize formal charges while ensuring the structure reflects the given total charge.
05

- Draw Lewis Structures

Draw the final Lewis structures with the formal charge indicated closest to the charged atom.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Valence Electrons
Valence electrons are the electrons in the outermost shell of an atom. These electrons are crucial because they participate in chemical bonding and reactions. When determining the Lewis structure for a molecule or ion, the first step is to count the valence electrons for each atom in the species.
To find the number of valence electrons:
  • Refer to the group number in the periodic table for each atom.
  • For charged species, add an electron for each negative charge or subtract one for each positive charge.
For example, in \text{-}\(\text{OH}\), oxygen has 6 valence electrons, and hydrogen has 1. Since there’s an additional negative charge, add 1 extra electron, making a total of 8. Understanding valence electrons helps determine how atoms will bond together in a molecule.
Skeletal Structures
The skeletal structure of a molecule shows how the atoms are connected by bonds. When drawing these structures:
  • Connect atoms using single lines representing single bonds.
  • The central atom is usually the least electronegative element, except for hydrogen.
By arranging the atoms logically and connecting them, you lay the foundation for adding electrons. In \text{-}\(\text{OH}\), connect the oxygen and hydrogen with a single line. Similarly, for \text{\(\text{BH}_4\)}, boron is the central atom connected to four hydrogens. Accurately drawing the skeletal structures sets you up correctly for distributing the electrons in the next steps.
Octet Rule
The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons. This rule helps predict the arrangement of electrons in a molecule:
  • Hydrogen follows the 'duet rule' and seeks 2 electrons.
  • Other atoms strive to complete an octet.
Distribute electrons to outer atoms first, ensuring each atom achieves its octet (or duet in the case of hydrogen). For instance, in \text{\(\text{NH}_4\)}, nitrogen needs 8 electrons around it, while each hydrogen needs only 2. The goal is to distribute the electrons such that all atoms satisfy the octet rule. This step is critical for determining the most stable Lewis structure.
Formal Charges
Formal charges are used to determine the most stable electron configuration in a molecule. The formal charge of an atom is calculated by:
  • Subtracting the number of electrons the atom 'owns' in the structure from the number of valence electrons it has.
The formula for calculating formal charge is: \text{Formal Charge} = \text{Valence Electrons} - \text{Non-Bonding Electrons} - ½(\text{Bonding Electrons}).
Adjust the electron distribution to minimize these charges while reflecting the total charge of the species. In \text{-}\(\text{Cl}\), chlorine should have 7 valence electrons, but because it gains an extra electron due to the negative charge, it has a formal charge of -1. Ensuring the formal charges are minimized often results in the most accurate representation of the molecule.
Electron Distribution
Electron distribution involves placing the remaining valence electrons in a way that satisfies atomic octets and minimizes formal charges. Follow these guidelines:
  • Place electrons around outer atoms first.
  • Then place any remaining electrons around the central atom.
  • Ensure that atoms like hydrogen only have 2 electrons.
For example:
  • In \text{-}\(\text{OH}\), after bonding, place the remaining electrons around the oxygen until its octet is complete.
Correct electron distribution helps in accurately drawing the final Lewis structure, showing how electrons interact and ensuring the stability of the molecule. This step is essential to complete the visual representation of the molecule or ion.

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Most popular questions from this chapter

Add electrons to complete the following Lewis structures. In each case, the charge is placed as close as possible to the charged atom. (a) \({ }^{+} \mathrm{CH}_{2}\) (b) \(-\mathrm{CH}_{2} \mathrm{CH}_{3}\) (c) \(\mathrm{HC}=\mathrm{CH}_{2}\) (d) \({ }^{+} \mathrm{OH}_{3}\) (e) \({ }^{-} \mathrm{OH}\) (f) \({ }^{+} \mathrm{NH}_{2}\) (g) \(-\mathrm{NH}_{2}\) (h) \(\mathrm{CH}_{3}-\mathrm{C} \equiv \mathrm{N}-\mathrm{H}\)

The bond between two fluorine atoms in \(\mathrm{F}_{2}\) can be viewed as resulting from the end-on interaction between the \(2 p_{y}\) orbital on each fluorine. Sketch the molecular orbitals produced through the interaction of two fluorine \(2 p_{y}\) atomic orbitals.

Write Lewis dot structures for the neutral diatomic molecules \(\mathrm{F}_{2}\) and \(\mathrm{N}_{2}\). In \(\mathrm{F}_{2}\), there is a single bond between the two atoms, but in \(\mathrm{N}_{2}\) there is a triple bond between the two atoms.

Draw another structure for nitromethane in which every atom is neutral. Hint: There are only single bonds in this structure.

Each of the following compounds has at least one multiple bond. Draw a Lewis structure for each molecule. Use lines to indicate electrons in bonds and dots to indicate nonbonding electrons. (a) \(\mathrm{F}_{2} \mathrm{CCF}_{2}\) *(b) \(\mathrm{H}_{3} \mathrm{CCN}\) (c) \(\mathrm{H}_{2} \mathrm{CO}\) (d) \(\mathrm{H}_{2} \mathrm{CCO}\) (e) \(\mathrm{H}_{2} \mathrm{CCHCHCH}_{2}\) (f) \(\mathrm{H}_{3} \mathrm{CNO}\) (g) \(\mathrm{H}_{3} \mathrm{COCOH}\)
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