Question

For each of the following bonds, draw a figure indicating the direction of the bond dipole, including which end of the bond is positive and which is negative. a. \(\mathrm{I}-\mathrm{I}\) b. \(\mathrm{I}-\mathrm{Br}\) c. \(\mathrm{I}-\mathrm{F}\) d. \(A t-I\)

Short answer

a. I-I: No bond dipole. b. I-Br: I (+) --> Br (-) c. I-F: I (+) --> F (-) d. At-I: At (+) --> I (-) In bonds b, c, and d, the bond dipole arrow points from the less electronegative atom (positive) to the more electronegative atom (negative).

Step-by-step solution

Identify the electronegativity of each atom involved in the bond

Before comparing electronegativities, we should know the approximated electronegativity values for the elements involved in the bond. Here, we use the Pauling electronegativity scale: - Iodine (I): 2.66 - Bromine (Br): 2.96 - Fluorine (F): 3.98 - Astatine (At): 2.2

Compare the electronegativity values and decide which atom is more electronegative

Now, compare the electronegativity values of the two atoms in each bond and determine the bond dipole direction. a. I-I: Since there is no difference in electronegativity (both have the same value of 2.66), there is no bond dipole. b. I-Br: Comparing the electronegativities, we have iodine (2.66) and bromine (2.96). Bromine is more electronegative, so the bond dipole points from iodine to bromine. c. I-F: Comparing the electronegativities, we have iodine (2.66) and fluorine (3.98), with fluorine being more electronegative. Hence, the bond dipole points from iodine to fluorine. d. At-I: Comparing the electronegativities, we have astatine (2.2) and iodine (2.66), with iodine being more electronegative. Thus, the bond dipole points from astatine to iodine.

Draw a figure indicating the direction of the bond dipole and the positive and negative ends

For each bond, draw a figure with an arrow pointing from the positive end (the less electronegative atom) to the negative end (the more electronegative atom). a. I-I: No bond dipole, so no figure needs to be drawn. b. I-Br: I (+) --> Br (-) The arrow points from Iodine to Bromine with a '+' sign on Iodine and a '-' sign on Bromine. c. I-F: I (+) --> F (-) The arrow points from Iodine to Fluorine with a '+' sign on Iodine and a '-' sign on Fluorine. d. At-I: At (+) --> I (-) The arrow points from Astatine to Iodine with a '+' sign on Astatine and a '-' sign on Iodine.

Key concepts

Electronegativity

Electronegativity is a fundamental concept to understand bond polarity.
It refers to the ability of an atom in a molecule to attract electrons towards itself.
In essence, it determines how strongly an atom will pull on shared electrons in a chemical bond. Electronegativity values are not measured directly, but rather assigned based on observed bonding behavior.
By comparing the electronegativity values of two bonded atoms, we can predict the behavior of the bond.

• If two atoms have identical electronegativity (as in the I-I bond), they share electrons equally, resulting in a nonpolar bond.
• When one atom is more electronegative, electrons are more attracted to it, making that end of the bond slightly more negative (as in the I-Br and I-F bonds).

This difference in electronegativity forms the basis for understanding bond dipoles.

Bond Dipole

The term **bond dipole** describes the separation of electrical charges along a bond.
A bond dipole occurs when two atoms in a bond have different electronegativities.
This results in a partial positive charge on the less electronegative atom, and a partial negative charge on the more electronegative atom.Bond dipoles can be visually represented using arrows.
The arrow points from the slightly positive atom (denoted as \(+\)) to the negative atom (denoted as \(-\)).
For example, in the I-F bond, the strong difference in electronegativity results in an arrow pointing from iodine to fluorine.Understanding bond dipoles helps predict molecule behavior.

• Molecules with strong bond dipoles tend to dissolve well in polar solvents like water.
• In larger molecules, individual bond dipoles can combine to produce an overall molecular dipole, impacting properties like boiling point and solubility.

Pauling Scale

The Pauling scale is the most commonly used measure of electronegativity.
It was developed by Linus Pauling, a key figure in chemistry, in the mid-20th century.
Pauling's scale assigns values to elements based on their chemical properties, with fluorine possessing the highest value of 3.98. With this scale, chemists can easily compare the electronegativities of different elements.
It provides an intuitive way to predict bond polarity and dipole moments.

• The greater the difference in electronegativity values between two atoms, the more polar the bond.
• When using the Pauling scale, electronegativity differences can also guide decisions in chemical reactions and during compound synthesis.

The simplicity and widespread acceptance of the scale make it a standard reference in understanding and teaching chemical bonding concepts.