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Chapter 4: Problem 87

What do each of the following sets of compounds/ions have in common with each other? a. \(\mathrm{XeCl}_{4}, \mathrm{XeCl}_{2}\) b. ICls, \(\operatorname{TeF}_{4}, \operatorname{ICl}_{3}, \mathrm{PCl}_{3}, \mathrm{SCl}_{2}, \mathrm{SeO}_{2}\)

Short Answer

Expert verified
Both sets of compounds have a central atom from Group 16 (chalcogens) in the periodic table. In the first set, the common feature is the presence of Xenon, a noble gas, and Chlorine. In the second set, the central atoms are Iodine, Tellurium, Phosphorus, Sulfur, or Selenium, and they bond with either Chlorine or Fluorine.

Step by step solution

01

Part A: Comparing \(\mathrm{XeCl}_{4}, \mathrm{XeCl}_{2}\)

To find the common characteristic between these two compounds, let's first examine their formulas. It is observed that both compounds are formed between Xenon (Xe), a noble gas, and Chlorine (Cl). The first molecule, \(\mathrm{XeCl}_{4}\), contains 4 Chlorine atoms bonded to the Xenon atom. This gives it a tetrahedral molecular geometry and a balanced electron distribution around the Xenon atom. The second molecule, \(\mathrm{XeCl}_{2}\), contains 2 Chlorine atoms bonded to the Xenon atom. This gives it a linear molecular geometry and a balanced electron distribution around the Xenon atom. To summarize, the common feature of these two compounds is that they are both compounds of Xenon (a noble gas) with Chlorine.
02

Part B: Comparing ICl5, \(\operatorname{TeF}_{4}, \operatorname{ICl}_{3}, \mathrm{PCl}_{3}, \mathrm{SCl}_{2}, \mathrm{SeO}_{2}\)

To find the common characteristic between these compounds, let's first examine their molecular formulas and structures. 1. ICl5 - Iodine hexachloride: 5 Chlorine atoms bonded to an Iodine atom, resulting in square pyramidal molecular geometry. 2. \(\operatorname{TeF}_{4}\) - Tellurium tetrafluoride: 4 Fluorine atoms bonded to a Tellurium atom, resulting in a see-saw molecular geometry. 3. \(\operatorname{ICl}_{3}\) - Iodine trichloride: 3 Chlorine atoms bonded to an Iodine atom, resulting in T-shaped molecular geometry. 4. \(\mathrm{PCl}_{3}\) - Phosphorus trichloride: 3 Chlorine atoms bonded to a Phosphorus atom, resulting in trigonal pyramidal molecular geometry. 5. \(\mathrm{SCl}_{2}\) - Sulfur dichloride: 2 Chlorine atoms bonded to a Sulfur atom, resulting in a bent molecular geometry. 6. \(\mathrm{SeO}_{2}\) - Selenium dioxide: 2 Oxygen atoms bonded to a Selenium atom, resulting in a bent molecular geometry. Notice that each compound consists of a central atom from Group 16 in the periodic table. The central atoms are I (Iodine), Te (Tellurium), Se (Selenium), or S (Sulfur). These elements belong to the same group, Group 16 (also known as the chalcogens). Therefore, the common feature of these compounds is that they are formed between Group 16 elements and either Chlorine or Fluorine.

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Key Concepts

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

Noble Gas Compounds
Understanding noble gas compounds is essential in the study of chemistry, as these compounds defy what was once considered a fixed rule—that noble gases are inert and do not form compounds due to their complete valence electron shells. However, some noble gases, particularly those heavier like Xenon, are known to form stable compounds.

Xenon, for instance, forms compounds like (XeCl_4) and (XeCl_2) with Chlorine. This is because Xenon has a sufficiently large atomic size that allows the outer electrons to be less tightly held, making it capable of chemically reacting under certain conditions. The resulting compounds exhibit interesting properties and molecular geometries. Understanding these unique compounds reinforces the concept that chemical reactivity can be quite nuanced and often defies simplistic rules.
Group 16 Elements
The elements of Group 16, also known as the chalcogens, include Oxygen, Sulfur, Selenium, Tellurium, and Polonium. These elements have diverse applications and play critical roles in biological systems and industrial processes. They all have six electrons in their outermost shell, which contributes to their typical -2 oxidation state.

Compounds like TeF_4, PCl_3, SCl_2, and SeO_2 show the versatility of Group 16 elements in forming diverse molecular structures with varied electron distributions and molecular geometries. The compounds often have covalent bonds and can exhibit different properties depending on the number and type of atoms they bond with. For instance, the bent molecular geometry of SCl_2 is a consequence of the electron pair geometry around the central sulfur atom.
Molecular Formulas
Molecular formulas convey critical information about a compound, including the types and numbers of atoms present. For example, the molecular formula for Iodine pentachloride is ICl_5, indicating one Iodine atom covalently bonded to five Chlorine atoms. Such a formula not only conveys the stoichiometric ratio but also helps in determining the compound's molecular weight and is essential in understanding the substance's stoichiometry in a chemical reaction. Moreover, the molecular formula serves as a gateway to predicting molecular geometry and the electron distribution around the central atom. Different molecular formulas denote varying shapes and bonding arrangements, which significantly impact the properties and reactivity of the compounds.
Electron Distribution
The electron distribution within a molecule ultimately dictates the molecule's shape, reactivity, and physical properties. By applying the principles of VSEPR (Valence Shell Electron Pair Repulsion) theory, we can predict the three-dimensional structure and electron distribution of molecules.

For instance, the linear geometry of (XeCl_2 ) is a result of the distribution of electron pairs around the Xenon atom, which minimizes repulsion between the pairs. Conversely, the see-saw molecular geometry of TeF_4 arises from the presence of four bonding pairs and one lone pair on the Tellurium atom. Analyzing electron distribution not only allows us to understand and predict molecular shapes but also elucidates the potential chemical behavior and intermolecular interactions that a compound can engage in.

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Most popular questions from this chapter

Describe the bonding in the \(\mathrm{O}_{3}\) molecule and the \(\mathrm{NO}_{2}^{-}\) ion, using the localized electron model. How would the molecular orbital model describe the \(\pi\) bonding in these two species?

Consider the molecular orbital electron configurations for \(\mathrm{N}_{2}\) \(\mathrm{N}_{2}^{+},\) and \(\mathrm{N}_{2}^{-} .\) For each compound or ion, fill in the table below with the correct number of electrons in each molecular orbital.

Which of the following are predicted by the molecular orbital model to be stable diatomic species? a. \(\mathrm{H}_{2}^{+}, \mathrm{H}_{2}, \mathrm{H}_{2}^{-}, \mathrm{H}_{2}^{2-}\) b. \(\mathrm{He}_{2}^{2+}, \mathrm{He}_{2}^{+}, \mathrm{He}_{2}\)

Give one example of a compound having a linear molecular structure that has an overall dipole moment (is polar) and one example that does not have an overall dipole moment (is nonpolar). Do the same for molecules that have trigonal planar and tetrahedral molecular structures.

Show how a hydrogen \(1 s\) atomic orbital and a fluorine \(2 p\) atomic orbital overlap to form bonding and antibonding molecular orbitals in the hydrogen fluoride molecule. Are these molecular orbitals \(\sigma\) or \(\pi\) molecular orbitals?
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