Question

Decide which bond in each pair is polar and explain your reasoning. Place \(\delta^{+}\)and \(\delta^{-}\)superscripts on the element symbols for each polar bond to indicate the direction of the polarity. (a) \(\mathrm{H}-\mathrm{F}\) and \(\mathrm{H}-\mathrm{H}\) (c) \(\mathrm{H}-\mathrm{H}\) and \(\mathrm{B}-\mathrm{H}\) (b) \(\mathrm{Cl}-\mathrm{I}\) and \(\mathrm{Cl}-\mathrm{Cl}\)

Short answer

In pair (a), the \(H-F\) bond is polar and marked as \(H\delta^{+}-F\delta^{-}\) and \(H-H\) bond is nonpolar. In pair (c), the \(H-H\) bond is nonpolar and \(B-H\) bond is polar and marked as \(B\delta^{+}-H\delta^{-}\). In pair (b), the \(Cl-I\) bond is polar and marked as \(Cl\delta^{+}-I\delta^{-}\) and \(Cl-Cl\) bond is nonpolar.

Step-by-step solution

Analyze Electronegativity Difference for each pair

Refer to a table of electronegativity values for the elements. The differences in electronegativity values for the atoms in each pair can determine the polarity. If the electronegativity difference is significantly great, then that bond is polar. If the electronegativity difference is zero or very small, the bond is nonpolar.

Apply Electronegativity Difference to Exercise Combinations

For pair (a) \(H-F\) and \(H-H\), \(H-F\) is polar because the electronegativity of Fluorine (3.98) is greater than Hydrogen (2.20) hence \(H\delta^{+}-F\delta^{-}\), while \(H-H\) is nonpolar as the atoms are identical. For pair (c) \(H-H\) and \(B-H\), \(H-H\) is still nonpolar, while \(B-H\) is polar since Boron (2.04) has lower electronegativity than Hydrogen, hence \(B\delta^{+}-H\delta^{-}\). For pair (b) \(Cl-I\) and \(Cl-Cl\), \(Cl-I\) is polar because Chlorine (3.16) has a higher electronegativity compared to Iodine (2.66), hence \(Cl\delta^{+}-I\delta^{-}\), while \(Cl-Cl\) is nonpolar as previously explained.

Conclude and Check

Having applied the electronegativity differences to each of the given pairs, we have determined the polarity for each required bond. Now check the reasoning to make sure the understanding of polar and non-polar bonds is correct, and that the direction of polarity has been indicated using \(\delta^{+}\)and \(\delta^{-}\)superscripts correctly according to electronegativity values.

Key concepts

Electronegativity

Electronegativity is a fundamental concept in chemistry that helps us understand the behavior of atoms in chemical bonds. It's essentially a measure of how strongly an atom attracts electrons in a bond. Each element has a specific electronegativity value, and these values can be found in electronegativity tables. The higher the value, the more an atom wants to "pull" electrons towards itself.

In the exercises, we looked at pairs like \(\mathrm{H}-\mathrm{F}\) and \(\mathrm{H}-\mathrm{H}\). We see that fluorine is much more electronegative than hydrogen, making the bond polar. This strong electronegativity of fluorine (3.98) shows its significant pull compared to hydrogen (2.20). By understanding these values, we can predict the behavior in a chemical bond and determine which side has more electron density.

Electronegativity differences help us gauge the bond polarity because they reveal which atom can "hog" more of the electron cloud when atoms bond. In essence, it's like a tug-of-war where the stronger "tugger" (electronegative element) pulls the electron cloud closer to itself, resulting in a polar bond.

Bond Polarity

Bond polarity emerges from the idea of unequal sharing of electrons between atoms in a chemical bond. When one atom in the bond has a higher electronegativity than the other, the electrons tend to spend more time closer to the more electronegative atom, creating a polar bond.

In our example, the bond \(\mathrm{H}-\mathrm{F}\) is polar because of the large difference in electronegativity between hydrogen and fluorine. We can represent this polarity using the notations \(\delta^{+}\) and \(\delta^{-}\). Here, \(\delta^{+}\) is assigned to hydrogen, indicating it is slightly positive since it has less electron density, and \(\delta^{-}\) is for fluorine because it attracts more electrons.

The same reasoning applies to \(\mathrm{B}-\mathrm{H}\), where hydrogen is more electronegative than boron, causing boron to be \(\delta^{+}\) and hydrogen \(\delta^{-}\). Bond polarity doesn't exist in \(\mathrm{H}-\mathrm{H}\) and \(\mathrm{Cl}-\mathrm{Cl}\) bonds because the atoms involved have the same electronegativity, leading to equal sharing of electrons.

Chemical Bonding

Chemical bonding is a fascinating force that holds atoms together within molecules. It involves the interaction of electrons, leading mainly to two types of bonding: covalent and ionic.

Covalent bonds occur when atoms share electrons. However, this sharing is not always equal, leading to polar and nonpolar bonds as we've seen earlier. The nature of the bond depends on the difference in electronegativity between the bonding atoms. In many cases, the sharing is unequal, resulting in a polar covalent bond, as with \(\mathrm{H}-\mathrm{F}\).

On the other hand, ionic bonds form when electrons are completely transferred from one atom to another, usually between atoms with a very high difference in electronegativity (typically more than 2.0). This transfer leads to the formation of charged ions, but in our examples, we're mostly looking at the polar covalent bonds where the transfer isn’t complete.

Understanding these concepts helps in predicting and explaining the properties of various compounds, like solubility and electrical conductivity, which are determined by the type of chemical bonding involved.