Question

A classmate of yours is convinced that he knows everything about electronegativity. (a) In the case of atoms \(X\) and \(Y\) having different electronegativities, he says, the diatomic molecule \(X-Y\) must be polar. Is your classmate correct? (b) Your classmate says that the farther the two atoms are apart in a bond, the larger the dipole moment will be. Is your classmate correct?

Short answer

(a) Your classmate is correct. Diatomic molecules with different electronegativities will generally be polar due to unequal sharing of electrons resulting in a net dipole moment. (b) Your classmate is not universally correct. While it may sometimes be true that a larger distance between atoms could lead to a larger dipole moment, this statement is not universally applicable. Many factors, such as the specific electronegativities of the atoms involved, could impact the dipole moment.

Step-by-step solution

Claim 1: Diatomic molecules with different electronegativities are polar

We need to verify if the general statement about diatomic molecules ("\(X-Y\)") with different electronegativities being polar is correct. A molecule is considered polar if it has a net dipole moment, meaning that there is a separation of positive and negative charges. This occurs when there is a difference in electronegativity between the atoms involved in the bond. In the case of a diatomic molecule, we have two atoms X and Y that are bonded together. Since the electronegativity values are different for X and Y, we have unequal sharing of electrons in the bond. Therefore, a dipole moment will be created, making the diatomic molecule polar. Thus, your classmate is correct in this claim.

Claim 2: Larger dipole moment when atoms are farther apart

The second claim that needs to be analyzed is whether the dipole moment is larger when the two atoms are farther apart in a bond. The dipole moment (μ), is represented by the equation μ = Q × d, where Q is the magnitude of the charge and d is the distance between the charged particles. In a bond, the distance between the nuclei of the two atoms is considered. Generally, as the distance between the two bonded atoms increases, the amount of sharing of electrons decreases, which can lead to a larger difference in the distribution of charge. Hence, with a larger separation of charges, the dipole moment could potentially be larger. However, we cannot be certain about this claim as many factors can come into play, such as the specific electronegativity values of the atoms in question, that could cause just the opposite result as well. Therefore, this statement is not universally correct. To conclude, Claim 1: Your classmate is correct. Diatomic molecules with different electronegativities will generally be polar due to unequal sharing of electrons resulting in a net dipole moment. Claim 2: Your classmate is not universally correct. While it may sometimes be true that a larger distance between atoms could lead to a larger dipole moment, this statement is not universally applicable. Many factors, such as the specific electronegativities of the atoms involved, could impact the dipole moment.

Key concepts

Polarity of Diatomic Molecules

When two atoms come together to form a diatomic molecule, one crucial factor that determines the nature of their bond is the concept of electronegativity. This term refers to the ability of an atom to attract shared electrons in a chemical bond. If atoms X and Y form a molecule and have different electronegativities, then the shared electrons are more likely to be found nearer to the more electronegative atom. This unequal distribution of electrons creates what is known as a polar bond.

A polar molecule has positive and negative ends, or poles, much like a magnet. For diatomic molecules, if the electronegativities are different, the molecule tends to be polar. A common example of this phenomenon is the hydrogen chloride (HCl) molecule. Chlorine is more electronegative than hydrogen, so the shared electrons are closer to the chlorine atom, leading to a partial negative charge there and a partial positive charge on the hydrogen atom. This separation of charges represents the polarity of the molecule, confirming what your classmate has claimed.

Dipole Moment

The dipole moment is a vector quantity that measures the separation of positive and negative charges within a molecule. It is represented by the Greek letter mu (μ). Mathematically, the dipole moment equation is given as \( \mu = Q \times d \) where \( Q \) represents the magnitude of the partial charges (assuming they are equal and opposite) and \( d \) is the distance between these charges. In essence, the dipole moment quantifies the polarity of a molecule.

Now, addressing the second part of the exercise, the idea that a larger distance between the atoms in a molecule leads to a larger dipole moment holds some truth. A greater separation means that the shared electrons are much more likely to be closer to one atom than the other, leading to a greater charge imbalance and a potentially larger \( \mu \) value. However, it is essential to understand that this is an oversimplification. Factors such as the precise electronegativity values and the type of chemical bond also play a significant role. For example, in molecules with relatively close electronegativity values, even a large distance may not significantly increase the dipole moment if the electron sharing remains relatively equal.

Chemical Bonding

The strength of the interaction between atoms X and Y in a diatomic molecule is defined by the chemical bond that holds them together. These bonds can be ionic, covalent, or metallic depending on the nature of the atoms involved and their electronegativities. For instance, when there is a significant difference in electronegativity, an ionic bond might be formed, where electrons are entirely transferred from one atom to another, as seen in sodium chloride (NaCl).

Covalent bonds, on the other hand, involve the sharing of electrons between atoms. The electronegativity difference between bonded atoms influences how evenly these electrons are shared. In the case of diatomic molecules, if one atom is more electronegative, it draws the shared electron pair closer to itself, creating a polar covalent bond. However, if the electronegativities are equal or very close, as seen in O₂ (oxygen) or N₂ (nitrogen), the electrons are shared equally, resulting in non-polar covalent bonds.

Understanding the type of chemical bond is crucial because it gives insight into the properties of the compound, such as boiling and melting points, solubility, and electrical conductivity. These properties highlight the importance of chemical bonding in understanding the behavior and characteristics of substances at a molecular level.