Question

Draw Lewis structures for the \(\mathrm{AsCl}_{4}^{+}\) and \(\mathrm{AsCl}_{6}^{-}\) ions. What type of reaction (acid-base, oxidation- reduction, or the like) is the following? $$ 2 \mathrm{AsCl}_{5}(g) \longrightarrow \mathrm{AsCl}_{4} \mathrm{AsCl}_{6}(s) $$

Short answer

The Lewis structures for the \(\mathrm{AsCl}_{4}^{+}\) and \(\mathrm{AsCl}_{6}^{-}\) ions are: \[ \hspace{4cm}\text{Cl}\overset{\ominus}{\text{As}}\left(\underset{\ominus}{\text{Cl}}\right)_3^{+1} \\ \hspace{4cm}\text{Cl}\overset{\ominus}{\text{As}}\left(\underset{\ominus}{\text{Cl}}\right)_5^{-1} \] The reaction of two \(\mathrm{AsCl}_5\) molecules to form an \(\mathrm{AsCl}_{4}^{+}\) ion and an \(\mathrm{AsCl}_{6}^{-}\) ion involves an electron transfer between the molecules, making it an oxidation-reduction (redox) reaction.

Step-by-step solution

Drawing the Lewis structure of \(\mathrm{AsCl}_{4}^{+}\) ion

To draw the Lewis structure of the \(\mathrm{AsCl}_{4}^{+}\) ion, first determine the total number of valence electrons in the molecule. Arsenic (As) has 5 valence electrons and each chlorine (Cl) atom has 7 valence electrons. The ion has a positive charge, so we must subtract 1 electron. Therefore, there are a total of 5 + 4(7) - 1 = 33 valence electrons. Place the As atom in the center and the four Cl atoms around it, forming single bonds between each Cl and the As atom. Distribute the remaining 29 valence electrons as lone pairs around the Cl atoms. Each Cl atom will have 3 lone pairs. The Lewis structure for the \(\mathrm{AsCl}_{4}^{+}\) ion is: \[ \hspace{4cm}\text{Cl}\overset{\ominus}{\text{As}}\left(\underset{\ominus}{\text{Cl}}\right)_3^{+1} \]

Drawing the Lewis structure of \(\mathrm{AsCl}_{6}^{-}\) ion

To draw the Lewis structure of the \(\mathrm{AsCl}_{6}^{-}\) ion, first determine the total number of valence electrons in the molecule. Arsenic (As) has 5 valence electrons and each chlorine (Cl) atom has 7 valence electrons. The ion has a negative charge, so we must add 1 electron. Therefore, there are a total of 5 + 6(7) + 1 = 47 valence electrons. Place the As atom in the center and the six Cl atoms around it, forming single bonds between each Cl and the As atom. Distribute the remaining 41 valence electrons as lone pairs around the Cl atoms. Each Cl atom will have 3 lone pairs. The Lewis structure for the \(\mathrm{AsCl}_{6}^{-}\) ion is: \[ \hspace{4cm}\text{Cl}\overset{\ominus}{\text{As}}\left(\underset{\ominus}{\text{Cl}}\right)_5^{-1} \]

Determine the type of reaction

To determine the type of reaction, analyze the changes occurring in the reaction. We have: $$ 2 \mathrm{AsCl}_5(g) \longrightarrow \mathrm{AsCl}_{4}^{+}(s) + \mathrm{AsCl}_{6}^{-}(s) $$ One \(\mathrm{AsCl}_5\) molecule is donating an electron to another \(\mathrm{AsCl}_5\) molecule. This leads to the formation of \(\mathrm{AsCl}_{4}^{+}\), which has one more Chlorine atom but is missing one of electron, and \(\mathrm{AsCl}_{6}^{-}\), which has one extra Chlorine atom and one extra electron. Since we observe an electron transfer from one molecule to another, this reaction is an oxidation-reduction (redox) reaction.

Key concepts

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. These electrons play a critical role in chemical bonding as they can be shared, gained, or lost during chemical reactions.
For example, when drawing Lewis structures, understanding the number of valence electrons helps indicate how atoms like arsenic and chlorine will bond.

• Arsenic (As) has 5 valence electrons.
• Chlorine (Cl) has 7 valence electrons.

In the exercise, determining the total number of valence electrons for each ion is essential. For \( \mathrm{AsCl}_{4}^{+} \), we calculate and adjust for the positive charge: 5 (from As) + 4(7) (from 4 Cl atoms) − 1 (due to the positive charge), resulting in 33 valence electrons. This understanding is crucial for placing electrons around each atom correctly.
Similarly, for the \( \mathrm{AsCl}_{6}^{-} \) ion, the addition of an extra electron due to the negative charge results in 47 valence electrons. Properly accounting for valence electrons ensures accuracy in Lewis structures and implies the possible formation of single bonds and lone pairs.

Oxidation-Reduction Reaction

Oxidation-reduction reactions, commonly referred to as redox reactions, involve the transfer of electrons between two species. One species undergoes oxidation by losing electrons, while the other undergoes reduction by gaining electrons.
In the reaction \(2 \mathrm{AsCl}_{5}(g) \longrightarrow \mathrm{AsCl}_{4} \mathrm{AsCl}_{6}(s)\), we see a classic example of a redox process, where electron transfer substantially alters the structure of the reactants:

• One molecule of \( \mathrm{AsCl}_5 \) loses an electron, forming \( \mathrm{AsCl}_{4}^{+} \), indicating oxidation as the oxidation number of As increases.
• The other molecule gains an electron, creating \( \mathrm{AsCl}_{6}^{-} \), showing reduction as the oxidation number of As decreases.

Recognizing changes in oxidation states (or charges of ions) during reactions can help identify redox reactions. Such electron transfers often result in the combination or separation of distinct chemical species, demonstrating the highly dynamic nature of redox chemistry.

Arsenic Chloride Ions

Arsenic chloride ions, such as \( \mathrm{AsCl}_{4}^{+} \) and \( \mathrm{AsCl}_{6}^{-} \), are specific molecules composed of arsenic bonded with chlorine. These ions can exhibit different charges, influencing their stability and reactivity. Here’s how these ions manifest:

• For \( \mathrm{AsCl}_{4}^{+} \), arsenic forms bonds with four chlorine atoms. It is a positively charged species because it loses an electron, relative to its bonding with neutral chlorine, leading to a reduced electron density.
• In \( \mathrm{AsCl}_{6}^{-} \), arsenic is bonded to six chlorine atoms. This ion is negatively charged, having gained an extra electron to accommodate the additional chloride ion.

These differences in structure and charge between \( \mathrm{AsCl}_{4}^{+} \) and \( \mathrm{AsCl}_{6}^{-} \) reflect their distinct roles in chemical processes. For example, in the provided reaction, the transformation from \( \mathrm{AsCl}_5 \) to these ions illustrates how ionic charges influence the types of products formed in redox reactions.
Understanding the formation and stability of such ions is essential for predicting chemical behavior and reactivity in more complex chemical systems.