Question

Which one of the following arrangements do not truely represent the property indicated against it? (a) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}<\mathrm{F}_{2}-\) Electronegativity (b) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}<\mathrm{F}_{2}-\) Bond energy (c) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}^{2}<\mathrm{F}_{2}-\) Electron affinity (d) \(\mathrm{Br}_{2}<\mathrm{Cl}_{2}<\mathrm{F}_{2}^{2}-\) oxidizing power

Short answer

Option B is incorrect for bond energy order: Cl₂ > Br₂ > F₂.

Step-by-step solution

Analyze Option A

Electronegativity refers to the ability of an atom to attract shared electrons. In the halogens, electronegativity decreases as you move down the group from fluorine to iodine. Therefore, the correct order for electronegativity is \( \text{F}_2 > \text{Cl}_2 > \text{Br}_2 \). Since option A states \( \text{Br}_2 < \text{Cl}_2 < \text{F}_2 \), which is correct, this is not the misrepresentation.

Analyze Option B

Bond energy refers to the amount of energy required to break a bond in a molecule. In diatomic halogens, bond energies generally decrease down the group, with the exception of \( \text{F}_2 \) due to its high reactivity and smaller bond length. So the usual order should be \( \text{Cl}_2 > \text{Br}_2 > \text{F}_2 \), not \( \text{Br}_2 < \text{Cl}_2 < \text{F}_2 \). Thus, Option B misrepresents bond energy.

Analyze Option C

Electron affinity is the amount of energy released when a neutral atom gains an electron. For halogens, electron affinity generally decreases down the group. The correct order should be \( \text{Cl}_2 > \text{F}_2 > \text{Br}_2 \) or \( \text{Cl}_2 > \text{Br}_2 > \text{F}_2 \). Option C misstates this order with \( \text{Br}_2 < \text{Cl}_2^2 < \text{F}_2 \), making it incorrect if interpreted off these norms.

Analyze Option D

Oxidizing power for halogens decreases down the group, meaning \( \text{F}_2 \) has the highest oxidizing power, followed by \( \text{Cl}_2 \) and then \( \text{Br}_2 \). The statement \( \text{Br}_2 < \text{Cl}_2 < \text{F}_2^2 \) correctly reflects the trend because \( \text{F}_2 \) is the strongest oxidizing agent.

Key concepts

Electronegativity

Electronegativity is a key concept in chemistry that measures an atom's tendency to attract a bonding pair of electrons. In the periodic table, electronegativity varies across periods and groups. For the halogens, which include fluorine (F), chlorine (Cl), and bromine (Br), electronegativity decreases as you move down the group.
The reason for this trend is due to the atomic structure. As you move down the group, the atoms have more electron shells. This increase in distance between the nucleus and the valence electrons reduces the nucleus's pull on shared electrons, thereby decreasing electronegativity.

• Fluorine is the most electronegative element, making it excellent at pulling electrons towards itself.
• Chlorine follows with a slightly lower electronegativity.
• Bromine has the lowest electronegativity among these three because its outer electrons are further from the nucleus compared to fluorine and chlorine.

Bond Energy

Bond energy refers to the amount of energy required to break one mole of a bond in a gaseous substance. It's a measure of the bond's strength, with a higher bond energy indicating a stronger bond. For diatomic molecules of halogens, such as *Fluorine ( F₂), chlorine ( Cl₂), and bromine ( Br₂)*, bond energies generally decrease as you move down the group.
Interestingly, fluorine’s bond energy is an exception. Despite being at the top of the group, F₂ has a surprisingly low bond energy. This anomaly arises due to the small size of the fluorine atom, which causes repulsions between non-bonding electrons when two fluorine atoms approach each other. This makes the F₂ bond relatively weak compared to Cl₂.

• Chlorine has the largest bond energy among the three, as the bond in Cl₂ experiences optimal overlap without significant repulsions.
• Bromine follows as it has a longer and therefore slightly weaker bond than chlorine.
• Fluorine, due to its electron repulsions, has the weakest bond energy among the three.

Electron Affinity

Electron affinity is defined as the energy change that occurs when an electron is added to a neutral atom, forming a negative ion. In simpler terms, it's a measure of an atom's ability to accept an electron. Halogens typically have high electron affinities because they only need one additional electron to obtain a full valence shell. However, trends can vary within the group.
Chlorine, contrary to what one might expect, has the highest electron affinity among F, Cl, and Br. This is because when an electron is added, the small size of the fluorine atom causes electron-electron repulsions that reduce its electron affinity compared to chlorine. As you continue down the group:

• Chlorine holds the highest electron affinity because its balance between size and nuclear charge makes it favorable for accepting an electron without significant repulsions.
• Despite being highly electronegative, fluorine's electron affinity is less than chlorine due to these repulsions.
• Bromine, with a larger atomic size, has a lower electron affinity compared to both chlorine and fluorine.

Oxidizing Power

Oxidizing power is a measure of a substance's ability to oxidize other substances, that is, to take away electrons. In other words, a stronger oxidizing agent will readily accept electrons and is more able to cause other substances to lose electrons.
For halogens, oxidizing power decreases as you move down the group from fluorine to iodine. The reason is tied to their electronegativity and ability to gain electrons:

• Fluorine is the most powerful oxidizing agent as it has a high electronegativity and a strong tendency to gain electrons.
• Chlorine follows with slightly less oxidizing power than fluorine.
• Bromine is less powerful than both fluorine and chlorine in oxidizing capacity.

The implication of this is that fluorine can more effectively oxidize other elements and compounds compared to chlorine and bromine.