Question

Write a Lewis structure that obeys the octet rule for each ion. Include resonance structures if necessary and assign formal charges to each atom. \begin{equation}\mathrm { a. }{ClO}_{3}^{-} \quad \text { b. } \mathrm{ClO}_{4}^{-} \quad \text { c. } \mathrm{NO}_{3}^{-} \quad \text { d. } \mathrm{NH}_{4}^{+}\end{equation}

Short answer

\({ClO}_{3}^{-}\): Central Cl with +2, three O each with -1; one O with a double bond. \({ClO}_{4}^{-}\): Central Cl with +1, all O with zero charge; no double bonds. \({NO}_{3}^{-}\): Central N with +1, one O with -1 in each resonance form; two single bonds and one double bond (three resonance structures). \({NH}_{4}^{+}\): N central atom with zero, all H with zero; no double bonds or formal charges.

Step-by-step solution

Determining Total Valence Electrons

For each ion, count the total number of valence electrons present. Take into account the charge of the ion (subtract an electron for a positive charge, add an electron for a negative charge).

Drawing the Skeletal Structure

Place the least electronegative atom in the center (except hydrogen, which always goes on the outside) and arrange the other atoms around it. Connect each outer atom to the central atom with a single bond.

Distributing Electrons to Obey the Octet Rule

Starting with the outer atoms, add electrons until they each fulfill the octet rule (eight electrons around each atom). Place any remaining electrons on the central atom.

Rechecking Electron Count and Making Double or Triple Bonds if Necessary

Ensure that the total number of electrons used equals the total number you calculated in step 1. If the central atom does not have an octet, form double or triple bonds as necessary by rearranging the outer electrons.

Drawing Resonance Structures

If there is more than one valid way to arrange double or triple bonds while maintaining the octet rule, draw all the possible resonance structures.

Assigning Formal Charges

Calculate the formal charge for each atom using the formula: Formal Charge = (Valence electrons) - (Non-bonding electrons) - (Bonding electrons/2). Assign the formal charge to each atom in the structure.

\({ClO}_{3}^{-}\) Lewis Structure

Chlorine (Cl) as the central atom, 26 valence electrons total, oxygen atoms (O) with six electrons each, add one extra electron for the charge. Distribute electrons to fulfill octet rule, form double bonds as necessary. One resonance structure. Formal charges: central Cl (+2), each O with single bond (-1).

\({ClO}_{4}^{-}\) Lewis Structure

Chlorine (Cl) as the central atom, 32 valence electrons total, oxygen atoms (O) with six electrons each, add one extra electron for the charge. Distribute electrons to fulfill octet rule, no double bonds needed. No resonance structures. Formal charge: central Cl (+1), oxygens with zero formal charge.

\({NO}_{3}^{-}\) Lewis Structure

Nitrogen (N) as the central atom, 24 valence electrons total, oxygen atoms (O) with six electrons each, add one extra electron for the charge. Distribute electrons to fulfill octet rule, form double bonds as necessary. Resonance structures possible. Formal charges: central N (+1), one O (-1) in each resonance form.

\({NH}_{4}^{+}\) Lewis Structure

Nitrogen (N) as the central atom, 8 valence electrons total (one subtracted for the charge), hydrogen atoms (H) with one electron each. No need for double bonds as hydrogen has a duet rule. No resonance structures. Formal charges: zero on all atoms.

Key concepts

Formal Charge Calculation

Understanding formal charge calculation is critical when creating Lewis structures. The formal charge is an indicator of the electron distribution within a molecule or ion. It is calculated using a simple formula:
\( \text{Formal Charge} = (\text{Valence electrons}) - (\text{Non-bonding electrons}) - \frac{1}{2}(\text{Bonding electrons}) \).
In this context, valence electrons are the electrons available for bonding in the atom's outermost shell. Non-bonding electrons are those not involved in bonding, while bonding electrons are shared in bonds and counted per bond.
For a well-formed Lewis structure, the sum of the formal charges should equal the overall charge of the molecule or ion. A good practice is to assign the least possible formal charge, ideally getting it close to zero, to ensure stability and accuracy of the structure.

Resonance Structures

Resonance structures provide a deeper understanding into the flexibility of electron arrangement in molecules and ions. When a molecule or ion has more than one valid Lewis structure, with electrons in different positions but the same arrangement of atoms, these are known as resonance structures. They are a way to represent delocalized electrons within certain molecules or ions that cannot be represented by a single Lewis structure.

Importance of Resonance

Resonance structures demonstrate how some molecules have a distribution of electrons that is more stable spread out over multiple atoms, rather than localized.
For example, in the provided solution, \( {NO}_{3}^{-} \) has possible resonance structures due to the ability to form different double bonds with oxygen atoms, reflecting electron delocalization. This contributes to the molecule's stability and is represented by arrows showing possible electron shifts.

Valence Electrons

Valence electrons play a fundamental role in chemical bonding, as they are the electrons found in an atom's outermost shell and are involved in forming bonds. The number of valence electrons an atom has determines its chemical properties and reactivity.
Atoms typically strive for a stable electron configuration, often represented by a full octet of valence electrons, following the 'octet rule'. This rule is a guiding principle for explaining the bonds that atoms form, but it does have exceptions.
In the exercise provided, counting valence electrons is the starting point for creating Lewis structures. For example, chlorine (Cl) has seven valence electrons, being in Group 17 of the periodic table. However, if it carries a charge, like in \( {ClO}_{3}^{-} \), you must adjust the electron count by adding one extra electron due to the negative charge, as it gains an extra electron.