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Mobile App Chapter 5: Problem 17
Draw a dot diagram for \(\mathrm{NO}^{+}\).
Short Answer
Expert verified
The dot diagram for the $\mathrm{NO}^{+}$ ion is as follows:
\[
\begin{array}{l}
\text{O} \\
= \\
\text{N}
\end{array}
\]
Step by step solution
01
Identify the Number of Valence Electrons
To begin, we will determine the number of valence electrons in a neutral molecule of NO. Nitrogen has 5 valence electrons (it is in group 15 on the periodic table), and Oxygen has 6 valence electrons (it is in group 16). Because this is a positive ion, however, it means there is one less electron. Therefore, the total number of valence electrons in the NO⁺ ion is:
\(5 (N) + 6 (O) - 1 (+) = 10\) valence electrons
We have a total of 10 valence electrons for our NO⁺ ion.
02
Draw the Skeleton Structure of NO⁺
Begin by drawing the skeleton structure of NO⁺: place the least electronegative atom (in this case, Nitrogen) in the center and the other atom (Oxygen) connected to it with a single bond.
```
O
|
N
```
Now, we have effectively put 2 of the 10 valence electrons to use in creating a single bond between Nitrogen and Oxygen. We have 8 electrons remaining to distribute.
03
Distribute Remaining Valence Electrons
Next, we need to complete the octet for the outermost shell (valence electrons) for each atom. Since there are 8 electrons left and Oxygen needs 6 more to complete its octet, we will place the remaining electrons in pairs around the Oxygen atom.
```
O
:|
N
```
Now, Oxygen has an octet, but Nitrogen has only 6 valence electrons around it, including the two forming the bond with Oxygen.
04
Adjust the Bonding
Since the octet rule is not fulfilled for Nitrogen (it lacks 2 electrons), we need to adjust the structure to ensure a complete octet. We can do this by sharing another pair of electrons from Oxygen with Nitrogen, forming a double bond.
```
O
=
N
```
Now, Nitrogen has a full octet (4 nonbonding electrons and 4 bonded electrons), and Oxygen has a full octet (2 nonbonding electron pairs and 2 bonded electron pairs). This is the final Lewis dot structure for the NO⁺ ion.
Result:
The dot diagram for NO⁺ is as follows:
```
O
=
N
```
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Valence Electrons
Understanding valence electrons is crucial when drawing Lewis dot structures, as these electrons play a key role in chemical bonding. Valence electrons are the electrons in the outermost energy level of an atom that are involved in forming bonds. Each atom aims to have a stable electron configuration, usually 8 valence electrons, known as an octet.
Elements in the same group of the periodic table have the same number of valence electrons, which is reflected in their chemical properties. For instance, all halogens have seven valence electrons and are highly reactive, often gaining one electron to achieve a stable octet. In contrast, noble gases already have a full octet, making them generally inert.
When determining the Lewis dot structure for an ion like \(\mathrm{NO}^{+}\), one must consider the charge when counting valence electrons. A positive charge indicates a loss of electrons, reducing the total available for bonding. In the \(\mathrm{NO}^{+}\) ion, the total number of valence electrons is calculated by adding the valence electrons from Nitrogen (5) and Oxygen (6) and subtracting one for the positive charge, resulting in ten valence electrons.
Elements in the same group of the periodic table have the same number of valence electrons, which is reflected in their chemical properties. For instance, all halogens have seven valence electrons and are highly reactive, often gaining one electron to achieve a stable octet. In contrast, noble gases already have a full octet, making them generally inert.
When determining the Lewis dot structure for an ion like \(\mathrm{NO}^{+}\), one must consider the charge when counting valence electrons. A positive charge indicates a loss of electrons, reducing the total available for bonding. In the \(\mathrm{NO}^{+}\) ion, the total number of valence electrons is calculated by adding the valence electrons from Nitrogen (5) and Oxygen (6) and subtracting one for the positive charge, resulting in ten valence electrons.
Octet Rule
The octet rule is a guiding principle in chemical bonding that refers to the tendency of atoms to prefer having eight electrons in their valence shell. This rule stems from the observation that atoms with a full octet resemble the electron configuration of noble gases, which are known for their stability and lack of reactivity.
When an atom doesn't naturally possess an octet, it will seek to achieve one through various types of bonding: sharing electrons (covalent bonding), losing electrons (ion formation), or gaining electrons (also ion formation). For example, in the Lewis dot structure of \(\mathrm{NO}^{+}\), Nitrogen initially has six valence electrons. It needs to form a bond by sharing electrons with Oxygen to fulfill the octet rule.
The structure is adjusted by creating a double bond between Nitrogen and Oxygen, allowing both to have an octet. This adherence to the octet rule is often but not universally accurate - it applies mainly to the second period elements. Some elements can have more (like Phosphorus or Sulfur) or fewer than eight electrons.
When an atom doesn't naturally possess an octet, it will seek to achieve one through various types of bonding: sharing electrons (covalent bonding), losing electrons (ion formation), or gaining electrons (also ion formation). For example, in the Lewis dot structure of \(\mathrm{NO}^{+}\), Nitrogen initially has six valence electrons. It needs to form a bond by sharing electrons with Oxygen to fulfill the octet rule.
The structure is adjusted by creating a double bond between Nitrogen and Oxygen, allowing both to have an octet. This adherence to the octet rule is often but not universally accurate - it applies mainly to the second period elements. Some elements can have more (like Phosphorus or Sulfur) or fewer than eight electrons.
Chemical Bonding
Chemical bonding is the force that holds atoms together to form molecules and compounds. The three main types are ionic, covalent, and metallic bonding.
In ionic bonding, electrons are transferred from one atom to another, resulting in the formation of positively and negatively charged ions. Covalent bonding involves the sharing of valence electrons between atoms, creating a stable balance by fulfilling the octet rule. Metallic bonding, on the other hand, is characterized by a 'sea' of shared electrons spread over a lattice of metal cations.
The Lewis dot structure primarily illustrates covalent bonding, as seen in \(\mathrm{NO}^{+}\), where electrons are shared between Nitrogen and Oxygen to obey the octet rule. This molecular ion, by sharing electrons more equally, exemplifies the fundamentals of covalent interactions, which are essential in the formation of many chemical compounds.
In ionic bonding, electrons are transferred from one atom to another, resulting in the formation of positively and negatively charged ions. Covalent bonding involves the sharing of valence electrons between atoms, creating a stable balance by fulfilling the octet rule. Metallic bonding, on the other hand, is characterized by a 'sea' of shared electrons spread over a lattice of metal cations.
The Lewis dot structure primarily illustrates covalent bonding, as seen in \(\mathrm{NO}^{+}\), where electrons are shared between Nitrogen and Oxygen to obey the octet rule. This molecular ion, by sharing electrons more equally, exemplifies the fundamentals of covalent interactions, which are essential in the formation of many chemical compounds.
