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Chapter 11: Problem 8

What does it mean to say that a bond is polar? Give two examples of molecules with polar bonds. Indicate in your examples the direction of the polarity.

Short Answer

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A polar bond is a covalent bond between two atoms where electrons are unequally shared due to differing electronegativity values, creating an electric dipole. Two examples of molecules with polar bonds are HCl and H2O. In HCl, the polarity points from H\((\delta^+)\) to Cl\((\delta^-)\), while in H2O, the polarity points from the H atoms (\(\delta^+\)) to the O atom (\(\delta^-\)).

Step by step solution

01

Definition of Polar Bond

A polar bond is a type of covalent bond between two atoms, where the electrons forming the bond are not equally shared by the two atoms due to differing electronegativity values. This results in a partially positive charge on the atom with lower electronegativity (and hence lesser electron-attracting power) and a partially negative charge on the atom with higher electronegativity. This creates an electric dipole, causing the bond to be polar.
02

Example 1: HCl (Hydrogen Chloride)

In a HCl molecule, the electron pair in the H-Cl bond is more attracted to the Cl atom due to its higher electronegativity value compared to H. As a result, the Cl atom acquires a partial negative charge (\(\delta^-\)), while the H atom acquires a partial positive charge (\(\delta^+\)). This leads to a polar bond, with the direction of polarity pointing from H to Cl, as shown below: H\((\delta^+)\)鈥擟l\((\delta^-)\)
03

Example 2: H2O (Water)

In a water molecule, the two O-H bonds are polar due to the difference in electronegativity values of O and H atoms. Since O is more electronegative than H, it attracts the bonded electron pairs more towards itself, resulting in a partial negative charge (\(\delta^-\)) on the O atom, and partial positive charges (\(\delta^+\)) on both H atoms. The directions of polarity point from H atoms to the O atom, as shown below: H 釐▅ O鈥擧 (\(\delta^+\)) (\(\delta^+\)) (\(\delta^-\))

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Most popular questions from this chapter

In each case, which of the following pairs of bonded elements forms the more polar bond? a. \(\mathrm{S}-\mathrm{F}\) or \(\mathrm{S}-\mathrm{Cl}\) b. \(\mathrm{N}-\mathrm{O}\) or \(\mathrm{P}-\mathrm{O}\) c. \(\mathrm{C}-\mathrm{H}\) or \(\mathrm{Si}-\mathrm{H}\)

Write a Lewis structure for each of the following simple molecules. Show all bonding valence electron pairs as lines and all nonbonding valence electron pairs as dots. a. \(\mathrm{H}_{2}\) b. HCl c. \(\mathrm{CF}_{4}\) d. \(\mathrm{C}_{2} \mathrm{F}_{6}\)

The ______ elements achieve an electron configuration analogous to the previous noble gas by losing electrons from their valence shells.

On the basis of their electron configurations, predict the formula of the simple binary ionic compound likely to form when the following pairs of elements react with each other. a. sodium, \(\mathrm{Na}\), and selenium, Se b. rubidium, Rb, and fluorine, F c. potassium, \(\mathrm{K}\), and tellurium, Te d. barium, Ba, and selenium, Se e. potassium, \(\mathrm{K}\), and astatine, \(\mathrm{At}\) f. francium, Fr, and chlorine, \(\mathrm{Cl}\)

For each of the following pairs, indicate which species is smaller. Explain your reasoning in terms of the electron structure of each species. a. \(L i^{+}\) or \(F^{-}\) b. \(\mathrm{Na}^{+}\) or \(\mathrm{Cl}^{-}\) c. \(\mathrm{Ca}^{2+}\) or \(\mathrm{Ca}\) d. \(\mathrm{Cs}^{+}\) or \(\mathrm{I}^{-}\)
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