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Chapter 8: Problem 36

List all the possible bonds that can occur between the elements \(\mathrm{P}, \mathrm{Cs}, \mathrm{O}\), and \(\mathrm{H}\). Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.

Short Answer

Expert verified
The possible bonds and their types between the elements P, Cs, O, and H are: - P-Cs: Polar covalent bond - P-O: Polar covalent bond - P-H: Non-polar covalent bond - Cs-O: Ionic bond - Cs-H: Polar covalent bond - O-H: Polar covalent bond

Step by step solution

01

1. P and Cs Bond

Computing the electronegativity difference between P and Cs: \(\Delta\mathrm{EN} = |2.19 - 0.79| = 1.4\) Since 0.4 < \(\Delta\mathrm{EN}\) < 1.7, the bond between P and Cs is polar covalent.
02

2. P and O Bond

Calculating the electronegativity difference between P and O: \(\Delta\mathrm{EN} = |2.19 - 3.44| = 1.25\) Since 0.4 < \(\Delta\mathrm{EN}\) < 1.7, the bond between P and O is polar covalent.
03

3. P and H Bond

Finding the electronegativity difference between P and H: \(\Delta\mathrm{EN} = |2.19 - 2.20| = 0.01\) Since \(\Delta\mathrm{EN}\) < 0.4, the bond between P and H is non-polar covalent.
04

4. Cs and O Bond

Computing the electronegativity difference between Cs and O: \(\Delta\mathrm{EN} = |0.79 - 3.44| = 2.65\) Since \(\Delta\mathrm{EN}\) > 1.7, the bond between Cs and O is ionic.
05

5. Cs and H Bond

Calculating the electronegativity difference between Cs and H: \(\Delta\mathrm{EN} = |0.79 - 2.20| = 1.41\) Since 0.4 < \(\Delta\mathrm{EN}\) < 1.7, the bond between Cs and H is polar covalent.
06

6. O and H Bond

Finding the electronegativity difference between O and H: \(\Delta\mathrm{EN} = |3.44 - 2.20| = 1.24\) Since 0.4 < \(\Delta\mathrm{EN}\) < 1.7, the bond between O and H is polar covalent. In summary: - P-Cs: Polar covalent bond - P-O: Polar covalent bond - P-H: Non-polar covalent bond - Cs-O: Ionic bond - Cs-H: Polar covalent bond - O-H: Polar covalent bond

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Most popular questions from this chapter

Indicate the bond polarity (show the partial positive and partial negative ends) in the following bonds. a. \(\mathrm{C}-\mathrm{O}\) d. \(\mathrm{Br}-\mathrm{Te}\) b. \(\mathrm{P}-\mathrm{H}\) e. \(\mathrm{Se}-\mathrm{S}\) c. \(\mathrm{H}-\mathrm{Cl}\)

Give an example of an ionic compound where both the anion and the cation are isoelectronic with each of the following noble gases. a. Ne c. \(\mathrm{Kr}\) b. \(\mathrm{Ar}\) d. Xe

Nitrous oxide \(\left(\mathrm{N}_{2} \mathrm{O}\right)\) has three possible Lewis structures: Given the following bond lengths, $$ \begin{array}{llll} \mathrm{N}-\mathrm{N} & 167 \mathrm{pm} & \mathrm{N}=\mathrm{O} & 115 \mathrm{pm} \\ \mathrm{N}=\mathrm{N} & 120 \mathrm{pm} & \mathrm{N}-\mathrm{O} & 147 \mathrm{pm} \\ \mathrm{N} \equiv \mathrm{N} & 110 \mathrm{pm} & & \end{array} $$ rationalize the observations that the \(\mathrm{N}-\mathrm{N}\) bond length in \(\mathrm{N}_{2} \mathrm{O}\) is \(112 \mathrm{pm}\) and that the \(\mathrm{N}-\mathrm{O}\) bond length is \(119 \mathrm{pm}\). Assign formal charges to the resonance structures for \(\mathrm{N}_{2} \mathrm{O}\). Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Combustion reactions of fossil fuels provide most of the energy needs of the world. Why are combustion reactions of fossil fuels so exothermic?

Given the following information: Heat of sublimation of \(\operatorname{Li}(s)=166 \mathrm{~kJ} / \mathrm{mol}\) Bond energy of \(\mathrm{HCl}=427 \mathrm{~kJ} / \mathrm{mol}\) Ionization energy of \(\operatorname{Li}(g)=520 . \mathrm{kJ} / \mathrm{mol}\) Electron affinity of \(\mathrm{Cl}(g)=-349 \mathrm{~kJ} / \mathrm{mol}\) Lattice energy of \(\operatorname{LiCl}(s)=-829 \mathrm{~kJ} / \mathrm{mol}\) Bond energy of \(\mathrm{H}_{2}=432 \mathrm{~kJ} / \mathrm{mol}\) Calculate the net change in energy for the following reaction: $$ 2 \mathrm{Li}(s)+2 \mathrm{HCl}(g) \longrightarrow 2 \mathrm{LiCl}(s)+\mathrm{H}_{2}(g) $$
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