Question

C-X bond is strongest in (a) \(\mathrm{CH}_{3} \mathrm{Br}\) (b) \(\mathrm{CH}_{3} \mathrm{Cl}\) (c) \(\mathrm{CH}_{3} \mathrm{I}\) (d) \(\mathrm{CH}_{3} \mathrm{~F}\)

Short answer

The C-X bond is strongest in \\( \mathrm{CH}_3 \mathrm{F} \\).

Step-by-step solution

Understanding Bond Strength

The bond strength in a molecule primarily depends on the bond length and the difference in electronegativities between the atoms forming the bond. Shorter bond lengths and larger differences in electronegativities typically result in stronger bonds.

Analyzing the C-X Bond

In the given options, we have C-X bonds where X is a halogen (Br, Cl, I, F). Bond strength generally decreases with increasing atomic size: as the atomic size of halogens increases, the bond length increases, making it weaker.

Examining Halogen Atomic Size

Among the halogens (F, Cl, Br, I), fluorine (F) is the smallest and iodine (I) is the largest. Shorter bonds (such as C-F) are stronger because the atoms are closer together, offering a stronger attraction.

Comparing Electronegativity

Fluorine is the most electronegative element, resulting in a highly polar and strong C-F bond due to the significant difference in electronegativity between C and F compared to other halogens.

Conclusion of Analysis

Considering both the bond length and the electronegativity difference, the C-F bond is the strongest among the options given. Thus, \( \mathrm{CH}_3 \mathrm{F} \) has the strongest C-X bond.

Key concepts

Electronegativity

Electronegativity is a measure of how strongly an atom attracts electrons in a bond. This concept is crucial in determining bond strength. When two atoms form a bond, the difference in their electronegativities affects the bond's characteristics.
Fluorine is the most electronegative element on the periodic table, surpassing other halogens like chlorine, bromine, and iodine. This means that when fluorine forms a bond with carbon, as in the case of a C-F bond, it pulls the shared electrons closer to itself more effectively than other halogens.

• A larger electronegativity difference between two atoms leads to a more polar bond.
• A highly polar bond is generally stronger and more stable.
• The C-F bond is exceptionally strong due to the significant electronegativity difference between carbon and fluorine.

Understanding electrical pull is vital for predicting bond strength. So, when analyzing these characteristics, taking into account the role of electronegativity is essential for assessing the behavior of bonds.

Atomic Size

Atomic size, or atomic radius, is the distance from the nucleus of an atom to its outermost electron cloud. This factor significantly influences bond strength. In general, as you move down a group in the periodic table, the atomic size increases due to the addition of electron shells.
For halogens, the order of increasing atomic size is: fluorine < chlorine < bromine < iodine.

• Fluorine, being the smallest halogen, allows for the closest approach to carbon, creating a short and strong bond.
• Bigger atoms like iodine result in longer bond lengths, which are generally weaker due to the decreased attraction between the bonded atoms.

The atomic size directly impacts the bond length. Shorter bonds, like C-F, are typically stronger because the atomic cores are closer, enabling a stronger interaction. Considering atomic size can provide a better understanding of why certain bonds exhibit greater strength and stability than others.

C-X Bond Analysis

Examining the C-X bond, where X represents a halogen such as fluorine, chlorine, bromine, or iodine, involves analyzing both bond length and electronegativity. Both factors are pivotal in determining bond strength.
In the case of C-F bonds, the combination of short bond length and substantial electronegativity difference makes them the strongest among C-X bonds. The C-F bond is the most robust due to fluorine's small atomic size and high electronegativity.

• Shorter bond lengths result in stronger bonds, as seen with C-F compared to other C-X bonds.
• The significant electronegativity difference enhances the bond polarity and strength.
• Analyzing both the atomic size and electronegativity offers a more comprehensive understanding of bond behavior.

This analysis aids in predicting the stability and reactivity of various bonds. It highlights the importance of considering multiple factors when assessing why some bonds are stronger than others.