Rank the members of each set of compounds in order of \(\mathrm{de}\) creasing ionic character of their bonds. Use partial charges to indicate the bond polarity of each: (a) \(\mathrm{PCl}_{3}, \mathrm{PBr}_{3}, \mathrm{PF}_{3}\) (b) \(\mathrm{BF}_{3}, \mathrm{NF}_{3}, \mathrm{CF}_{4}\) (c) \(\mathrm{SeF}_{4}, \mathrm{TeF}_{4}, \mathrm{BrF}_{3}\)

Electronegativity is a key concept in chemistry, and it helps us understand how atoms attract shared electrons in a bond.
Electronegativity values are based on a scale developed by Linus Pauling, where fluorine has the highest value of 3.98.
Other elements have lower electronegativities:

• Phosphorus (P): 2.19
• Chlorine (Cl): 3.16
• Bromine (Br): 2.96
• Boron (B): 2.04
• Nitrogen (N): 3.04
• Carbon (C): 2.55
• Selenium (Se): 2.55
• Tellurium (Te): 2.10

When comparing these values, the larger the difference between the atoms, the more the shared electrons are pulled towards the more electronegative atom. This concept is essential in predicting bond behavior and understanding the ionic character of compounds.

When we talk about ionic character, we are looking at how much a bond between two atoms resembles an ionic bond.
In an ionic bond, electrons are completely transferred from one atom to another, usually between metals and non-metals.
In covalent bonds, electrons are shared. However, in polar covalent bonds, this sharing is unequal because of differences in electronegativity.

To determine the ionic character of a bond, we compare the electronegativity values of the atoms involved.
A greater difference in electronegativity means a higher ionic character, indicating a stronger pull of electrons towards the more electronegative atom.

For example, in PF3, the difference in electronegativity between P (2.19) and F (3.98) is 1.79.
This is larger compared to PCl3, where the difference between P (2.19) and Cl (3.16) is 0.97. Thus, PF3 has higher ionic character.

Bond polarity is a result of differences in electronegativity between two bonding atoms. Polarity in a bond means that there are slight charges (partial charges) on both atoms.
The more electronegative atom becomes slightly negative (δ-), while the less electronegative atom becomes slightly positive (δ+).

A bond with a large difference in electronegativity will be more polar. Take BF3 as an example:

• Boron (B) electronegativity: 2.04
• Fluorine (F) electronegativity: 3.98

The difference in this case is 1.94, which means the B-F bond is quite polar, with B bearing a δ+ and F bearing a δ-.
This polarity can affect the compound's properties, such as solubility and melting point.

Partial charges (denoted as δ+ and δ-) are a way of showing the unequal distribution of electrons in polar covalent bonds.
They are not full charges like in ionic compounds but indicate fractional electron density shifts.
These partial charges arise because the more electronegative atom attracts shared electrons more strongly.

Consider NF3:
Here, nitrogen (N) has an electronegativity of 3.04, and fluorine (F) has 3.98.
This results in a partial negative charge on fluorine (Fδ-) and a partial positive charge on nitrogen (Nδ+).
Similarly, CF4 has carbon (C) with 2.55 and fluorine (F) also at 3.98, creating similar partial charges but less intense due to a slightly lower electronegativity difference (1.43).
This principle helps predict molecule behavior and interactions, such as hydrogen bonding and molecule orientation in electric fields.