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Chapter 5: Problem 95

Draw a dot diagram for the nitrate \(\left(\mathrm{NO}_{3}^{-}\right)\) ion.

Short Answer

Expert verified
The Lewis structure for the nitrate ion (NO3-) can be drawn as: O \ N - O / O- In this structure, nitrogen is in the center and connected to each of the three oxygen atoms with single bonds. Two oxygen atoms have single bonds and six non-bonding electrons, while one oxygen atom has a double bond and four non-bonding electrons, giving that oxygen a formal charge of -1. The overall charge of the ion is -1, and the nitrate ion exhibits resonance.

Step by step solution

01

Determine the Number of Valence Electrons

We need to find the total number of valence electrons for the nitrogen, the three oxygen atoms, and the extra electron from the negative charge: - Nitrogen (N) has 5 valence electrons - Oxygen (O) has 6 valence electrons - The ion has a negative charge, so we need to add 1 electron Total valence electrons: \(5 + 3 \times 6 + 1 = 24\)
02

Connect the Atoms with Single Bonds

Place the nitrogen atom in the center, as it has the lowest electronegativity. Connect the nitrogen atom to each of the three oxygen atoms using single bonds. This uses up 2 electrons per bond, for a total of 6 electrons used.
03

Complete the Octets

Now we want to complete the octet for each oxygen atom by adding pairs of electrons around them. We have 18 remaining electrons to distribute (24 total valence electrons minus the initial 6 used for single bonds). Each oxygen atom needs 6 more electrons (2 pairs) to complete its octet. We need to assign these electrons to the three oxygen atoms. Assign 2 pairs of electrons to each oxygen atom. This uses up all the remaining 18 electrons.
04

Check the Octets and Formal Charges

We check to make sure that the octet rule is satisfied for all atoms and that the formal charges are correct: - Nitrogen (N) is surrounded by 3 single bonds, so it has a formal charge of 5 - 3 = +2 - Each oxygen (O) is surrounded by 1 single bond and 6 non-bonding electrons, so it has a formal charge of 6 - (1 + 6) = 0 - The formal charge of the ion is +2 - 1 = +1 (adding one electron to satisfy the negative charge of the ion) We still need to add one electron (the extra negative charge) to one of the oxygen atoms.
05

Add the Extra Electron

Add the extra electron to one of the oxygen atoms, completing the octet for that oxygen atom. This gives that particular oxygen atom a formal charge of -1, and the overall charge of the ion becomes -1.
06

Draw the Final Dot Diagram

Now we can draw the final dot diagram (Lewis structure) for the nitrate ion (NO3-): O \ N - O / O- Keep in mind that the double bond can be on any of the three oxygen atoms, making the nitrate ion a resonance structure.

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Key Concepts

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

Valence Electrons
Valence electrons play a crucial role in chemical bonding as they are the outermost electrons of an atom and participate in forming bonds with other atoms. For the nitrate ion (l(NO3^-)), understanding valence electrons is the first step in drawing its dot diagram. Nitrogen (N) has 5 valence electrons, and each oxygen (O) has 6. Since the ion carries an additional negative charge, we include an extra electron. The sum of which defines the electron count that governs how atoms within the ion will bond and arrange themselves.
Lewis Structure
The Lewis structure is a visual representation that depicts the arrangement of atoms and the distribution of valence electrons around them. When creating a Lewis structure for the nitrate ion, we place the nitrogen in the center and connect it to each oxygen atom through single bonds. Afterward, by adding electron pairs around each oxygen atom, we attempt to complete their octet, which ultimately results in the clear visual arrangement of electrons that adheres to the octet rule.
Octet Rule
The octet rule is a guiding principle in chemistry that states atoms are most stable when they have eight electrons in their valence shell, mimicking noble gases' electron configuration. In context to the nitrate ion, this rule is used to determine the placement of remaining electrons after forming bonds, specifically aiming to complete the octets around the oxygen atoms while ensuring nitrogen鈥檚 formal charge and the ion鈥檚 charge are balanced.
Formal Charge
Calculating the formal charge for each atom in a molecule helps to determine the most stable electron structure. The formal charge is the difference between the valence electrons an atom owns in a molecule versus its isolated state. In the nitrate ion, nitrogen has more bonds than valence electrons, leading to a positive formal charge. Oxygen atoms typically have a zero formal charge when they complete their octet without forming a double bond. However, when accommodating the nitrate鈥檚 negative charge, one oxygen ends up with an extra electron, resulting in a negative formal charge, thus making the entire ion's charge balanced.
Resonance Structure
Resonance structures occur when there are multiple ways to depict a molecule's electron structure with Lewis diagrams, without changing the atoms' positions. In the nitrate ion, the double bond can be with any of the three oxygen atoms, and thus, the ion has three equivalent resonance structures. This depicts the delocalization of electrons and provides a more accurate representation of the distribution of electrons within the molecule.

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Most popular questions from this chapter

For any representative element, what is the relationship between the element's Romannumeral group number and the number of dots in the Lewis dot diagram?

Iodine atoms in \(\mathrm{I}_{2}\) should have a (a) \(1^{-}\) charge (b) \(\delta^{-}\) charge (c) \(\delta^{+}\) charge (d) No charge

Chlorine (Cl), neon (Ne), and helium (He) all exist as gases, but only one of them is diatomic. Which is it, and why is it diatomic while the others are monatomic?

The oxide ion is \(\mathrm{O}^{2-}\). How does this ion differ from the peroxide ion, \(\mathrm{O}_{2}^{2-} ?\) Draw dot diagrams for both.

The greater the electronegativity difference between two bonded atoms, the greater the percent character of the bond.
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