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Chapter 9: Problem 70

Write Lewis structures for the reaction $$ \mathrm{AlCl}_{3}+\mathrm{Cl}^{-} \longrightarrow \mathrm{AlCl}_{4}^{-} $$ What kind of bond joins Al and \(\mathrm{Cl}\) in the product?

Short Answer

Expert verified
The Lewis structures show a central Al atom bonded to three (or four in AlCl4-) peripheral Cl atoms, with each Cl having three lone pairs of electrons. Al and Cl form ionic bonds in the product AlCl4-.

Step by step solution

01

Identify the reactants and the product

The reactants in this chemical reaction are AlCl3 and Cl-, while the product is AlCl4-.
02

Write Lewis Structures for the reactants

The Lewis structure for AlCl3 shows a central Al atom bonded to three peripheral Cl atoms. Each Cl atom has three lone pairs of electrons and shares a pair with Al. The Cl- ion is represented as a Cl atom surrounded by four lone pairs of electrons which include the extra electron, indicating a negative charge.
03

Write Lewis Structure for the product

In the product, AlCl4-, the central Al atom is bonded to four peripheral Cl atoms. Each Cl atom has three lone pairs of electrons and shares a pair with Al. Also, the molecule has a negative charge because it has gained an electron.
04

Identify the type of bond

The type of bond formed between Al and Cl is ionic (more specifically polar covalent) as Aluminum (metal) tends to donate electrons and Chlorine (non-metal) tends to accept electrons, forming cations and anions.

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Most popular questions from this chapter

Write Lewis structures for the following ions: (a) \(\mathrm{O}_{2}^{2-},\) (b) \(\mathrm{C}_{2}^{2-},\) (c) \(\mathrm{NO}^{+}\), (d) \(\mathrm{NH}_{4}^{+}\). Show formal charges.

The amide ion, \(\mathrm{NH}_{2}^{-},\) is a Bronsted base. Represent the reaction between the amide ion and water.

Because fluorine has seven valence electrons \(\left(2 s^{2} 2 p^{5}\right),\) seven covalent bonds in principle could form around the atom. Such a compound might be \(\mathrm{FH}_{7}\) or \(\mathrm{FCl}_{7}\). These compounds have never been prepared. Why?

A new allotrope of oxygen, \(\mathrm{O}_{4}\), has been reported. The exact structure of \(\mathrm{O}_{4}\) is unknown, but the \(\operatorname{sim}-\) plest possible structure would be a four-member ring consisting of oxygen-oxygen single bonds. The report speculated that the \(\mathrm{O}_{4}\) molecule might be useful as a fuel "because it packs a lot of oxygen in a small space, so it might be even more energy-dense than the liquefied ordinary oxygen used in rocket fuel." (a) Draw a Lewis structure for \(\mathrm{O}_{4}\) and write a balanced chemical equation for the reaction between ethane, \(\mathrm{C}_{2} \mathrm{H}_{6}(g),\) and \(\mathrm{O}_{4}(g)\) to give carbon dioxide and water vapor. (b) Estimate \(\Delta H^{\circ}\) for the reaction. (c) Write a chemical equation illustrating the standard enthalpy of formation of \(\mathrm{O}_{4}(g)\) and estimate \(\Delta H_{\mathrm{f}}^{\circ}\) (d) Assuming the oxygen allotropes are in excess, which will release more energy when reacted with ethane (or any other fuel): \(\mathrm{O}_{2}(g)\) or \(\mathrm{O}_{4}(g) ?\) Explain using your answers to parts (a)-(c).

Write three resonance structures for (a) the cyanate ion \(\left(\mathrm{NCO}^{-}\right)\) and \((\mathrm{b})\) the isocyanate ion \(\left(\mathrm{CNO}^{-}\right) .\) In each case, rank the resonance structures in order of increasing importance.
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