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Chapter 8: Problem 148

List the bonds $\mathrm{P}-\mathrm{Cl}, \mathrm{P}-\mathrm{F}, \mathrm{O}-\mathrm{F},\( and \)\mathrm{Si}-\mathrm{F}$ from least polar to most polar.

Short Answer

Expert verified
The order of bonds from least polar to most polar is \(\mathrm{O}-\mathrm{F}, \mathrm{P}-\mathrm{Cl}, \mathrm{P}-\mathrm{F},\) and \(\mathrm{Si}-\mathrm{F}\).

Step by step solution

01

Determine the electronegativity values of the elements involved in the bonds

: Consult a periodic table or reference material to find the electronegativity values of the elements involved in the bonds: P (Phosphorus) - Electronegativity value: 2.19 Cl (Chlorine) - Electronegativity value: 3.16 F (Fluorine) - Electronegativity value: 3.98 O (Oxygen) - Electronegativity value: 3.44 Si (Silicon) - Electronegativity value: 1.90
02

Calculate the electronegativity differences for each bond

: Find the difference in electronegativity values for each bond. The absolute value of the difference will be used to compare polarities: P-Cl: |2.19 - 3.16| = 0.97 P-F: |2.19 - 3.98| = 1.79 O-F: |3.44 - 3.98| = 0.54 Si-F: |1.90 - 3.98| = 2.08
03

Order the bonds from least to most polar

: Now that we have the electronegativity differences for each bond, we can order them from least polar to most polar based on these differences: 1. O-F (0.54) 2. P-Cl (0.97) 3. P-F (1.79) 4. Si-F (2.08) So, the order of bonds from least polar to most polar is \(\mathrm{O}-\mathrm{F}, \mathrm{P}-\mathrm{Cl}, \mathrm{P}-\mathrm{F},\) and \(\mathrm{Si}-\mathrm{F}\).

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Most popular questions from this chapter

The most common type of exception to the octet rule are compounds or ions with central atoms having more than eight electrons around them. PF_. \(\mathrm{PF}_{5}, \mathrm{SF}_{4}, \mathrm{ClF}_{3}\) and \(\mathrm{Br}_{3}^{-}\) are examples of this type of exception. Draw the Lewis structure for these compounds or ions. Which elements, when they have to, can have more than eight electrons around them? How is this rationalized?

Predict the molecular structure (including bond angles) for each of the following. (See Exercises 115 and 116.) a. \(\mathrm{ICl}_{5}\) b. \(\mathrm{XeCl}_{4}\) c. \(\mathrm{SeCl}_{6}\)

Use the following standard enthalpies of formation to estimate the \(\mathrm{N}-\mathrm{H}\) bond energy in ammonia: $\mathrm{N}(g), 472.7 \mathrm{kJ} / \mathrm{mol} ; \mathrm{H}(g),\( \)216.0 \mathrm{kJ} / \mathrm{mol} ; \mathrm{NH}_{3}(g),-46.1 \mathrm{kJ} / \mathrm{mol}$ . Compare your value to the one in Table \(8.5 .\)

Carbon and sulfur form compounds with each other with the formulas \(\mathrm{CS}_{2}\) and \(\mathrm{C}_{3} \mathrm{S}_{2}\) . Draw a Lewis structure for each compound that has a formal charge of zero for all atoms in the structure.

Arrange the atoms and/or ions in the following groups in order of decreasing size. a. \(\mathrm{O}, \mathrm{O}^{-}, \mathrm{O}^{2-}\) b. \(\mathrm{Fe}^{2+}, \mathrm{Ni}^{2+}, \mathrm{Zn}^{2+}\) c. \(\mathrm{Ca}^{2+}, \mathrm{K}^{+}, \mathrm{Cl}^{-}\)
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