Question

Describe the type of bonding that exists in the \(\mathrm{F}_{2}(g)\) molecule. How does this type of bonding differ from that found in the \(\mathrm{HF}(g)\) molecule? How is it similar?

Short answer

In the \(\mathrm{F}_{2}(g)\) molecule, non-polar covalent bonding occurs, as both fluorine atoms have the same electronegativity and equally share the electron pair. On the other hand, the \(\mathrm{HF}(g)\) molecule exhibits polar covalent bonding due to the higher electronegativity of fluorine compared to hydrogen, leading to an unequal sharing of the electron pair and creating a dipole moment. However, both molecules have covalent bonds as they share electron pairs to achieve a stable state.

Step-by-step solution

Observing the F_2 molecule

The given molecule \(\mathrm{F}_{2}(g)\) consists of two fluorine (F) atoms joined together. In this case, since both atoms are of the same element (fluorine), the type of bonding between the atoms would be covalent, as they share a pair of electrons to achieve a full octet and reach stability. Covalent bonds occur typically between non-metallic elements, which is the case with fluorine.

Observing the HF molecule

Now, for the other molecule, \(\mathrm{HF}(g)\), it consists of a hydrogen (H) and a fluorine (F) atom combined. To identify the nature of bonding, we have to consider the electronegativity difference between the two elements involved. Fluorine is highly electronegative, while hydrogen is of comparatively lower electronegativity. Due to this electronegativity difference, fluorine attracts the shared pair of electrons more towards itself, creating a polar covalent bond in the \(\mathrm{HF}\) molecule.

Comparing F_2 and HF bonding types

Comparing the type of bonding found in the \(\mathrm{F}_{2}(g)\) and \(\mathrm{HF}(g)\) molecules, we can conclude the following: 1. Similarity: - Both \(\mathrm{F}_{2}\) and \(\mathrm{HF}\) exhibit covalent bonding as they share a pair of electrons to achieve a stable state (an octet in the case of fluorine and the H-atom reaches a stable duet in the case of hydrogen). 2. Differences: - The bonding in \(\mathrm{F}_{2}\) is non-polar covalent bonding, as both fluorine atoms have the same electronegativity, resulting in equal sharing of the electron pair. - In \(\mathrm{HF}\), the bonding is polar covalent, as fluorine attracts the shared electron pair more towards itself, due to its higher electronegativity compared to hydrogen. This creates a dipole moment in the molecule, resulting in the polar nature of the bond. In conclusion, both \(\mathrm{F}_{2}(g)\) and [\(\mathrm{HF}(g)\) exhibit covalent bonding, with \(\mathrm{F}_{2}(g)\) being non-polar and \(\mathrm{HF}(g)\) being polar, due to the differences in electronegativity between the elements involved in each molecule.

Key concepts

Covalent Bond

A covalent bond is a chemical bond that forms when two atoms share a pair of electrons to achieve stability. This type of bond primarily happens between non-metallic elements. A great example of covalent bonding is seen in the gaseous fluorine molecule (\(mathrm{F}_{2}(g)\)). In this molecule, two fluorine atoms come together and each shares one electron with the other. This sharing results in each atom having a complete outer shell, achieving the much-desired stable configuration. Since both atoms involved belong to the same element—fluorine in this case—there is no electronegativity difference, and the electrons are shared equally.Covalent bonds are integral in forming the vast majority of organic molecules. They provide the necessary stability and structure that characterize compounds containing carbon, nitrogen, oxygen, and other non-metals. Here are some fast facts about covalent bonds:

• They occur mostly between non-metal elements.
• They involve sharing electron pairs.
• They help atoms achieve a full outer electron shell.

Polar Covalent Bond

A polar covalent bond arises when two atoms share electrons, but the electrons are not shared equally. This situation typically happens between atoms of different electronegativities. For instance, in the molecule hydrogen fluoride (\(mathrm{HF}(g)\)), the electronegativity difference between hydrogen and fluorine causes the shared electron pair to be pulled more towards the fluorine atom. This creates an unequal sharing where fluorine becomes slightly negatively charged while hydrogen takes on a slightly positive charge.Polar covalent bonds lead to the formation of dipole moments, which can impact the molecule's physical properties. These bonds are crucial in explaining the behavior of molecules in different environments. Characteristics of polar covalent bonds:

• They occur between atoms with different electronegativities.
• They result in partial charges within the molecule.
• They influence how molecules interact with one another.

In summary, polar covalent bonds are essential for creating molecules with unique chemical behaviors and interaction patterns. They are especially important in biological systems and many industrial chemical processes.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons. It is a fundamental concept that helps understand the behavior of atoms in a molecule. Essentially, it determines how electrons are distributed in a chemical bond.The electronegativity scale was developed by Linus Pauling, where fluorine was assigned the highest electronegativity of 3.98. This ranking makes fluorine incredibly effective at attracting electrons, which is evident in polar covalent bonds such as those found in the \(mathrm{HF}(g)\) molecule.Here are some critical points to remember about electronegativity:

• Atoms with high electronegativity tend to pull electrons towards themselves.
• Electronegativity can predict bond type: differences lead to polarity in bonds.
• It assists in predicting molecule shapes, bond energies, and other properties.

Understanding electronegativity helps in anticipating how molecules will interact under various conditions and explains why specific reactions occur the way they do. It acts as a guiding principle for predicting molecule reactivity and properties, making it a cornerstone concept in chemistry.