Problem 21 - Chemistry in 30 days

Question

Which one of the following arrangements not truly represent the property inlin ated ayainst it? (a) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}<\mathrm{F}_{2}\) : Electro negativity (b) \(\mathrm{Br}_{2}<\mathrm{F}_{2}<\mathrm{Cl}_{4}\) : I lectron affinity (c) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}<\mathrm{F}_{2}:\) Bond energy (d) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}<\mathrm{I}_{1}\), Onidirimg priwer

Step-by-Step Solution

Answer

Option (c) does not correctly represent the property it claims, while option (b) misrepresents electron affinity trends.

Step 1: Understand Electronegativity

Electronegativity refers to the ability of an atom to attract shared electrons in a bond. In general, electronegativity increases across a period in the periodic table and decreases down a group. Thus, for halogens: \(\mathrm{F}_{2} > \mathrm{Cl}_{2} > \mathrm{Br}_{2}\). Option (a) correctly aligns with this trend.

Step 2: Examine Electron Affinity

Electron affinity measures the energy change when an electron is added to a neutral atom. Typically, more negative electron affinity means a higher tendency to gain electrons. A common trend is \(\mathrm{F}_{2} > \mathrm{Cl}_{2} > \mathrm{Br}_{2}\), contrary to option (b). For halogens, \(\mathrm{Cl}_{2}\) usually has a more negative electron affinity than \(\mathrm{F}_{2}\).

Step 3: Check Bond Energy

Bond energy for diatomic molecules involves understanding the strength required to break the bond. Normally, across the halogens: \(\mathrm{Cl}_{2} > \mathrm{F}_{2} > \mathrm{Br}_{2}\). Hence, option (c) does not follow the general trend for bond energy.

Step 4: Evaluate Oxidizing Power

Oxidizing power is the ability of an element to oxidize another substance. The general trend for halogens is: \(\mathrm{F}_{2} > \mathrm{Cl}_{2} > \mathrm{Br}_{2}\). Thus, according to option (d), it aligns with the typical oxidizing capability trend.

Key Concepts

Understanding Electronegativity

Electronegativity is a fundamental property that defines how atoms attract and hold onto electrons in a chemical bond. This peculiar ability of an atom is crucial as it influences the chemical reactivity and bonding nature of molecules.

Across the periodic table, electronegativity follows a specific trend:

• Increases from left to right across a period.
• Decreases from top to bottom down a group.

In the context of halogens, fluorine ({F}) is known for having the highest electronegativity due to its small size and high effective nuclear charge. Therefore, in sequences involving halogens like {F_2}, {Cl_2}, and {Br_2}, we expect fluorine to be the most electronegative, followed by chlorine and then bromine. This reflects the inherent tendency of fluorine to attract shared electrons more effectively compared to chlorine and bromine.

Understanding electronegativity helps predict how strongly atoms can pull electrons towards themselves, affecting bond polarity and various physical and chemical properties of compounds.

Exploring Electron Affinity

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom. It tells us how readily an atom accepts additional electrons. When an atom has a high electron affinity, it releases more energy upon gaining an electron, indicating a higher tendency to accept electrons.

Typically, halogens exhibit high electron affinities because gaining an electron allows them to achieve a noble gas configuration:

• Fluorine, being very electronegative, might seem to have the highest electron affinity. However, due to its small atomic size, repulsions between electrons can reduce this ability slightly.
• Chlorine, despite being less electronegative than fluorine, often shows a higher, more negative electron affinity as it strikes a balance between electronegativity and atomic size.

Hence, for halogens like chlorine and fluorine, chlorine occasionally showcases a higher electron affinity compared to fluorine, contrary to expected trends based solely on electronegativity.

This concept of electron affinity is crucial in understanding reactivity patterns, especially among elements in the same group.

Deciphering Bond Energy

Bond energy is the measure of the strength of a chemical bond, specifically the amount of energy needed to break one mole of bonds in gaseous molecules. It reflects the bond's stability and strength.

Generally, counterintuitive trends can occur in halogens like:

• Chlorine molecules ({Cl_2}) having a surprisingly higher bond energy than fluorine ({F_2}), contrary to what might be expected from their atomic size and electronegativity.
• This is owed to the closer distance between the fluorine atoms, leading to increased repulsion, which decreases bond strength when compared to chlorine.

Thus, while we'd expect the smallest halogen, fluorine, to have a stronger bond, its atom-atom repulsions cause a weaker bond compared to chlorine.

Understanding bond energy is vital for predicting molecular stability and reactivity, as well as the energy changes involved in chemical reactions.

Unpacking Oxidizing Power

The oxidizing power of a substance refers to its ability to oxidize other substances, particularly through gaining electrons or causing electron loss in another entity. It is crucial as it defines how substances interact in redox reactions.

For halogens, the oxidizing power is often aligned with their electronegativity:

• Fluorine ({F_2}) stands out as the strongest oxidizing agent due to its high electronegativity and capacity to attract electrons.
• It surpasses even chlorine ({Cl_2}) and bromine ({Br_2}) in its ability to oxidize other elements.

This capacity of halogens decreases as we go down the group, making fluorine the most potent oxidizer among them. Its exceptional ability to pull electrons away from other elements is what drives its high oxidizing strength.

Comprehending oxidizing power helps explain many chemical reactions involving halogens, especially their role in oxidizing other compounds.